MIEMCAL 


COLLEGE  OF  PHARMACY 


GENERAL  CHEMISTRY 
FOR  COLLEGES  ' 


BY 

ALEXANDER    SMITH 

PROFKSSOK  OF  CHEMISTRY  AND  DIRECTOR  OF  GENERAL  AND  PHYSICAL  CHEMISTRY 
IN  THE  UNIVERSITY  OF  CHICAGO 


..,»  .  «*«  r=Q*!e<*'©  o 
California 


NEW  YORK 

THE   CENTURY  CO, 

1908 


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Stanbope 

P.   H.   OIL80N   COMPANY 
BOSTON.   O.  8.  A. 


PREFACE. 


THE  present  work  differs  from  the  Author's  "Introduction  to 
General  Inorganic  Chemistry"  in  being  intended  for  pupils  who 
can  devote  less  time  to  the  study  of  the  science,  and  whose  needs 
can  be  satisfied  by  a  less  extensive  course.  It  resembles  the 
larger  work  in  the  arrangement  of  the  contents  and  in  the  general 
method  of  treatment.  The  matter,  and  particularly  the  theoreti- 
cal matter,  however,  has  been  simplified  and  has  been  confined 
strictly  to  the  most  fundamental  topics.  Such  parts  of  the  theory 
as  are  thus  given,  are  presented  with  the  same  fullness  as  before, 
and  are  illustrated  and  applied  with  all  the  persistence  needed  to 
insure  full  apprehension  and,  ultimately,  spontaneous  employment 
by  the  student.  Such  parts  as  could  not  be  treated  in  this  way, 
within  the  limits  set  by  the  plan  of  the  book,  have  been  omitted. 
Methods  materially  different  from  those  used  in  the  "Introduction" 
have  been  employed  in  presenting  many  topics.  Conspicuous 
differences  of  this  kind  will  be  noted  particularly  in  the  treatment 
of  combining  proportions,  formulae  and  equations,  molecular  and 
atomic  weights,  chemical  equilibrium,  ionic  substances  and  their 
interactions,  and  the  theory  of  precipitation. 

The  writer  desires  to  express  his  profound  gratitude  to  the  many 
chemists  who  have  made  valuable  criticisms  and  suggestions. 
Most  of  these  comments  applied  to  the  "Introduction  to  General 
Inorganic  Chemistry,"  but  many  of  them  have  been  used  in 
preparing  this  work  (General  Chemistry  for  Colleges),  and  all  will 
be  considered  in  the  second  edition  of  the  larger  book. 

For  critical  reading  of  the  whole  of  the  proofs  of  the  present 
work,  the  writer  desires  especially  to  thank  Messrs.  A.  T.  McLeod 
and  Alan  W.  C.  Menzies  of  the  University  of  Chicago.  Other  cor- 
rections and  suggestions  will  be  gladly  received  by  the  author. 

ALEXANDER  SMITH. 
CHICAGO,  April,  1908. 


CONTENTS.' 


CHAPTERS  PAaE3 

I.     INTRODUCTORY  I 1 

The  Scientific  Method  —  Three  Illustrative  Chemical  Phe- 
nomena—  Two  Characteristics  of  Chemical  Phenomena. 

II.     INTRODUCTORY  II 12 

Conservation  of  Mass  —  Energy  —  Applications  of  the  Con- 
ception of  Energy  in  Chemistry  —  Chemical  Activity  and  its 
Cause  —  Of  Elements  and  of  Simple  and  Compound  Sub- 
stances —  Some  of  the  Fundamental  Ideas  Used  by  Chemists 
and  the  Corresponding  Terms  —  Methods  of  Work  and  Ob- 
servation in  Chemistry. 

HI.     INTRODUCTORY  III 29 

The  Law  of  Definite  Proportions  —  The  Measurement  of 
Combining  Weights  —  The  Law  of  Multiple  Proportions  — 
The  Law  of  Combining  Weights  —  Equivalents  —  Atomic 
Weights. 

IV.     INTRODUCTORY  IV 39 

Symbols,  Formulae,  and  Equations  —  Units  of  Measurement 
in  Chemical  Work  —  Calculations  in  Chemistry. 

V.    OXYGEN      45 

Preparation  of  Oxygen  —  Physical  Properties  of  Oxygen  — 
Chemical  Properties  of  Oxygen  —  The  Making  of  Equations 
Again  —  Oxides  —  Combustion  —  Oxidation  —  Means  of 
Altering  the  Speed  of  a  Given  Chemical  Action:  By  Change 
of  Temperature  —  Rapid  Self-sustaining  Chemical  Action  and 
Means  of  Initiating  It  —  Other  Means  of  Altering  the  Speed 
of  a  Given  Chemical  Change  :  By  Catalysis  —  Thermo- 
chemistry. 

*  The  titles  of  most  of  the  descriptive  paragraphs  have  been  omitted  from  this  tahle 
because  it  is  easier  to  find  matters  of  this  nature  by  consulting  the  index.  The  titles  of 
all  the  theoretical  paragraphs,  however,  have  been  included. 

vii 


viii  CONTENTS 

CHAPTERS  PAGES 

VI.     THE  MEASUREMENT  OF  QUANTITY  IN  GASES 58 

The  Measurement  of  the  Pressure  of  a  Gas  —  Boyle's  Law  — 
The  Correction  of  the  Volume  of  a  Gas  for  Temperature  — 
Mixed  Gases  —  Densities  of  Gases. 


VII.    HYDROGEN 63 

Acids  —  Preparation  of  Hydrogen  —  Tests  —  Preparation  of 
Simple  Substances  —  Displacement  —  Valence  —  Physical 
Properties  of  Hydrogen  —  Diffusion  —  Chemical  Properties 
of  Hydrogen  —  Specific  Chemical  Properties  —  The  Speed 
of  Chemical  Actions  :  A  Means  of  Measuring  Activity. 


VIII.     WATER 77 

Natural  Waters  —  Physical  Properties  of  Water  — •  Ice  — 
Steam  and  Aqueous  Tension  —  Chemical  Properties  of 
Water  —  Hydrates  —  Composition  of  Water  —  Gay-Lussac's 
Law  of  Combining  Volumes. 


IX.     THE  KINETIC-MOLECULAR  HYPOTHESIS 86 

Kinetic-Molecular  Hypothesis  Applied  to  Gases  —  Kinetic 
Hypothesis  Applied  to  Liquids  —  Equilibrium  —  Kinetic 
Hypothesis  Applied  to  Solids  —  Crystal  Forms. 


X.    SOLUTION 96 

General  Properties  of  Solutions  —  The  Scope  of  the  Word  — 
Limits  of  Solubility  —  Recognition  and  Measurement  of  Solu- 
bility —  Terminology  —  Solution  one  of  the  Physical  States 
of  Aggregation  of  Matter  —  Kinetic-Molecular  Hypothesis 
Applied  to  the  State  of  Solution  —  Kinetic-Molecular 
Hypothesis  Applied  to  the  Process  of  Solution  —  Independent 
Solubility  —  Two  Immiscible  Solvents:  Law  of  Partition  — 
Influence  of  Temperature  on  Solubility  —  Equilibrium  in  a 
Saturated  Solution. 


XI.    CHLORINE  AND  HYDROGEN  CHLORIDE 108 

Chlorine  —  Chemical  Relations  of  the  Element  —  Hydrogen 
Chloride  —  Interaction  of  Acids  and  Chlorides  —  The 
Kinetic  Hypothesis  Applied  to  the  Interaction  of  Sulphuric 
Acid  and  Salt  —  Classification  of  Chemical  Interactions  and 
Exercises  Thereon. 


CONTENTS  ix 

CHAPTERS  PAGES 

XII.    MOLECULAR  WEIGHTS  AND  ATOMIC  WEIGHTS 125 

Meaning  of  Avogadro's  Hypothesis  —  Molecular  Weights  — 
Atomic  Weights  —  Advantages  of  Atomic  Weights  over 
Equivalents — Dulong  and  Petit 's  Law — Molecular  Formulae 
of  Compounds  —  Molecular  Formulae  of  Simple  Substances 
Applications  :  Interactions  between  Gases,  Molecular  Equa- 
tions, Arithmetical  Problems,  Cases  of  Dissociation,  Finding 
the  Atomic  Weight  of  a  New  Element  —  Replies  to  Ques- 
tions about  Difficulties  —  Exercises. 

XIII.  THE  ATOMIC  HYPOTHESIS 151 

XIV.  THE  HALOGEN  FAMILY 157 

The  Chemical  Relations  of  Elements  —  The  Chemical  Rela- 
tions of  the  Halogens  • —  Bromine  —  Partial  Equations,  a 
Plan  for  Making  Complex  Equations  —  Hydrogen  Bromide 

—  Iodine  — -    Hydrogen    Iodide    —  Fluorine    —  Hydrogen 
Fluoride  — •  The  Halogens  as  a  Family  —  Compounds  of  the 
Halogens  with  Each  Other. 

XV.     CHEMICAL  EQUILIBRIUM 174 

Reversible  Actions  —  Kinetic  Explanation  —  Chemical 
Equilibrium  and  its  Characteristics  —  The  Influence  of  Tem- 
perature —  The  Influence  of  Concentration  —  Formulation 
of  the  Law  of  Molecular  Concentration  —  Applications  :  The 
Forward  Action,  Homogeneous  and  Inhomogeneous  Systems 

—  Applications:    The    Reverse    Action.     Displacement    of 
Equilibria. 

XVI.    OXIDES  AND  OXYGEN  ACIDS  OF  THE  HALOGENS 186 

Compounds  of  Chlorine  Containing  Oxygen  —  Nomenclature 
of  Acids  and  Salts  —  Salts  and  Double  Decomposition  — 
Hypochlorites  and  Hypochlorous  Acid  —  Hypochlorous 
Anhydride  —  Hypochlorous  Acid  as  an  Oxidizing  Agent  : 
Bleaching  —  Simultaneous,  Independent  Chemical  Changes 
in  the  Same  Substance  —  Chlorates  —  Chloric  Acid  —  Chlo- 
rine Dioxide:  Chlorous  Acid  —  Perchlorates,  Perchloric  Acid 
and  Perchloric  Anhydride  —  Oxygen  Acids  of  Bromine  — 
Oxides  and  Oxygen  Acids  of  Iodine  —  Chemical  Relations, 

XVII.    DISSOCIATION  IN  SOLUTION 201 

Some  Characteristic  Properties  of  Acids,  Bases,  and  Salts, 
Shown  in  Aqueous  Solution  —  Freezing  Points  of  Solutions 
—  Laws  of  Freezing-point  Depression  —  Freezing-Points 
and  Dissociation  in  Solution  —  Applications:  The  Constitu- 
tion of  Solutions  of  Acids,  Bases,  and  Salts. 


x  CONTENTS 

CHAPTERS  PAGES 

XVIII.     OZONE  AND  HYDROGEN  PEROXIDE 209 

XIX.     IONIZATION 214 

Introductory  —  Non-Electrolytes  —  Chemical  Changes 
Taking  Place  in  Electrolysis:  At  the  Electrodes  —  Ionic 
Migration  —  The  Hypothesis  of  Ions  —  Application  to 
the  Explanation  of  Electrolysis  —  Difficulties  Presented  by 
this  Hypothesis  —  Re'sume'  and  Nomenclature  —  Applica- 
tions :  Ionic  Equilibrium  —  Applications  to  the  Interpre- 
tation of  Conductivity  Measurements  —  Constitution  of 
Solutions  of  lonogens  :  Fractions  Ionized  —  Relation  of 
lonization  to  Chemical  Activity. 

XX.    IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS 231 

The  classes  of  lonogens:  Mixed  lonogens  and  Double  Salts 
—  The  Kinds  of  Ionic  Substances  Furnished  by  lono- 
gens —  Ionic  Equilibrium  with  a  Single  lonogen  —  Dis- 
placement of  Ionic  Equlibria  —  Applications  :  Double 
Decomposition  in  Solution  —  Precipitation  —  Neutraliza- 
tion —  Acidimetry  and  Alkalimetry  —  Indicators  - 
Displacement  :  Electromotive  Series  —  Non-Ionic  Modes 
of  Forming  lonogens. 

XXI.    SULPHUR  AND  HYDROGEN  SULPHIDE 248 

Sulphur  —  Hydrogen  Sulphide  —  Sulphides  —  Polysul- 
phides  —  The  Chemical  Relations  of  the  Element  Sulphur. 

XXII.     THE  OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR 256 

Sulphur  Dioxide  —  Sulphur  Troxide  —  Oxygen  Acids  of 
Sulphur  —  Sulphuric  Acid  —  Sulphates  —  Sulphurous 
Acid  —  Sulphites  —  Thiosulphuric  Acid  —  Persulphuric 
Acid  —  Sulphur  Monochloride. 

XXIII.  SELENIUM  AND  TELLURIUM  :  THE  PERIODIC  SYSTEM   ....     269 

Selenium  —  Tellurium  —  The  Chemical  Relations  of  the 
Sulphur  Family  —  Metallic  and  Non-metallic  Elements  — 
Classification  by  Atomic  Weights  —  Mendelejeff' s  Scheme 
—  General  Relations  in  the  System  —  Applications  of  the 
Periodic  System. 

XXIV.  NITROGEN  AND  AMMONIA 279 

Nitrogen  —  Ammonia  —  Ammonium  Compounds  — Hy- 
drazine  —  Hydrazoic  Acid  —  Halogen  Compounds  of 
Nitrogen. 


CONTENTS 


XI 


HAPTERS  PAGES 

XXV.     THE  ATMOSPHERE.     THE  HELIUM  FAMILY 286 

Components  of  the  Atmosphere  —  Air  a  Mixture  —  Lique- 
faction of  Gases  —  Argon  —  Helium  —  Neon,  Krypton 
and  Xenon. 

XXVI.     OXIDES  AND  OXYGEN  ACIDS  OF  NITROGEN 292 

Nitric  Acid  —  Nitric  Oxide  —  Nitrogen  Tetroxide  —  Oxi- 
dizing Actions  of  Nitric  Acid  —  Nitrous  Acid  — 
Nitrites  —  Nitrous  Oxide  —  Explosives  —  Active  ("  nas- 
cent") Hydrogen. 

XXVII.     PHOSPHORUS 303 

Phosphorus  —  The  Electric  Furnace  —  Phosphine  —  Hal- 
ides  of  Phosphorus  —  Oxides  of  Phosphorus  —  The  Phos- 
phoric Acids  and  Their  Salts  —  Phosphorous  Acid  —  Com- 
parison of  Phosphorus  with  Nitrogen  and  with  Sulphur. 

[XVIII.     CARBON  AND  THE  OXIDES  OP  CARBON 316 

Carbon  —  Calcium  Carbide — Carbon  Dioxide  and  Carbonic 
Acid  —  Carbonates  —  Role  of  Chlorophyll-bearing  Plants 
in  Storing  Energy  —  Carbon  Monoxide  —  Carbon  Disul- 
phide. 

XXIX.     SOME  CARBON  COMPOUNDS 327 

The    Hydrocarbons  —  Petroleum  —  Methane  —  Ethylene 

—  Acetylene  —  Formic  Acid  —  Acetic  Acid  —  Oxalic  Acid 

—  Carbohydrates  and   Fermentation  —  Alcohols,    Esters, 
and  Soap  —  Cyanogen  and  Derivatives. 

XXX.     FLAME  AND  ILLUMINANTS 337 

XXXI.     SILICON  AND  BORON 343 

Silicon  —  Silicon  Hydride  —  Carbide  of  Silicon  —  Silicon 
Tetrachloride  and  Tetrafluoride  — Hydrofluosilicic  Acid  — 
Silicon     Dioxide  —  Silicic     Acid  —  Silicates  —  Boron  — 
Hydrides  and  Halides  of  Boron  —  Boric  Acid  —  Boron 
Trioxide  —  Borates. 


XXXII.    THE  BASE-FORMING  ELEMENTS 351 

Physical  Properties  of  the  Metals  —  General  Chemical  Rela- 
tions of  the  Metallic  Elements  —  Occurrence  of  the  Metallic 
Elements  in  Nature  —  Methods  of  Extraction  from  the 
Ores  —  Compounds  of  the  Metals:  Oxides  and  Hydrox- 
ides — Compounds  of  the  Metals :  Salts. 


Xll 


CONTENTS 


CHAPTERS 
XXXIII. 


XXXIV. 


XXXV. 


XXXVI. 


XXXVII. 


PAGES 
THE  METALLIC  ELEMENTS  OF  THE  ALKALIES:  POTASSIUM 

AND  AMMONIUM 360 

The  Chemical  Relations  of  the  Metallic  Elements  of 
the  Alkalies  —  Potassium  —  the  Hydride  —  Chloride — 
Other  Halides  —  Hydroxide  —  Oxides  —  Chlorate  and 
lodate  —  Nitrate  —  Carbonate  —  Cyanide  —  Sulphate 
and  Bisulphate  —  Sulphides  —  Properties  of  Potassium- 
Ion  :  Analytical  Reactions  —  Rubidium  and  Caesium  — - 
Ammonium  —  Salts  of  Ammonium  —  Ammonium  Amal- 
gam —  Ammonium-Ion  :  Analytical  Reactions. 

SODIUM  AND  LITHIUM.    IONIC   EQUILIBRIUM   CONSIDERED 

QUANTITATIVELY 373 

SODIUM  :  Sodium  Chloride  —  Hydroxide  and  Oxides  — 
Nitrate  and  Nitrite  —  Carbonate  —  Bicarbonate  — 
Other  Salts  —  Properties  of  Sodium-Ion  :  Analytical 
Reactions  —  Lithium. 

IONIC  EQUILIBRIUM  CONSIDERED  QUANTITATIVELY:  Ex- 
cess of  One  Ion  —  Formulation  and  Quantitative  Treat- 
ment of  the  Case  of  Excess  of  one  Ion  —  Special  Case  of 
Saturated  Solutions  —  Illustration  of  the  Principle  of 
Ion-Product  Constancy  —  Other  Illustrations. 

THE  METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS.  .  .  388 
The  Chemical  Relations  of  the  Elements  —  Calcium  and 
its  Compounds  —  Hard  Water  —  Mortar  and  Cement  — 
Theory  of  Precipitation  —  Rule  for  Solution  of  Sub- 
stances —  Application  of  the  Rule  for  Solution  to  the 
Solution  of  Insoluble  Substances  —  Interaction  of  Insol- 
uble Salts  with  Acids  Resulting  in  Solution  of  the  Salt  — 
— Precipitation  of  Insoluble  Salts  in  Presence  of  Acids  — 
Glass  —  Calcium-Ion  :  Analytical  Reactions  —  Stron- 
tium and  its  Compounds  —  Barium  and  its  Compounds 
—  Analytical  Reactions  of  the  Calcium  Family. 

COPPER,  SILVER,  GOLD 408 

The  Chemical  Relations  of  the  Copper  Family  —  Copper 
and  its  Compounds  —  The  Solution  of  Insoluble  Salts 
when  Complex  Ions  are  Formed  —  Silver  and  its  Com- 
pounds —  Electro-plating  —  Photography  —  Gold  and 
its  Compounds. 


GLUCINUM,  MAGNESIUM,  ZINC,  CADMIUM,  MERCURY.    THE 
RECOGNITION  OF  CATIONS  IN  QUALITATIVE  ANALYSIS 


427 


CONTEXTS 


XI 11 


IHAPTKRS 

iXXVIII. 


ALUMINIUM  AND  THE  METALS  OF  THE  EARTHS 


PAGES 
442 


The  Rare  Elements  of  This  Family  —  Aluminium  and 
its  Compounds  —  Dyeing  :  Mordanting  —  Kaolin  and 
Clay  :  Earthenware  and  Porcelain. 

XXXIX.     GERMANIUM,  TIN,  LEAD 451 

XL.     ARSENIC,  ANTIMONY,  BISMUTH 462 

XLI.     THE  CHROMIUM  FAMILY.     RADIUM 474 

The  Chemical  Relations  of  the  Family  —  Chromium  — 
Derivatives  of  Chromic  Acid  —  Chromic  and  Chromous 
Compounds  —  Analytical  Reactions  —  Molybdenum  — 
Tungsten  —  Uranium  —  Radium  —  The  Discovery  of 
the  Element  —  Properties  of  Radium  Compounds  —  The 
Decay  of  an  Element. 

XLII.     MANGANESE 487 

XLIII.     IRON,  COBALT,  NICKEL 494 

XLIV.     THE  PLATINUM  METALS 508 

Ruthenium  and   Osmium  —  Rhodium   and   Indium  — 
Palladium  and  Platinum. 


GENERAL    CHEMISTRY    FOR   COLLEGES 


GENEKAL  CHEMISTRY  FOR  COLLEGES 

CHAPTER  I 
INTRODUCTORY   I 

As  human  knowledge  has  increased,  it  has  become  necessary  to 
subdivide  into  groups  the  growing  accumulation  of  truths,  and  to 
do  this  in  a  somewhat  arbitrary  fashion.  Each  such  group  is 
called  a  science,  and  includes  a  more  or  less  distinct  body  of  knowl- 
edge. That  the  boundaries  of  these  groups  do  not  exist  in  the 
subject-matter  itself  is  seen  at  once  in  our  own  treatment  of  them. 
Thus,  for  convenience,  we  include  in  the  science  of  physiology  the 
study  of  the  way  in  which  the  parts  of  both  plants  and  animals 
perform  their  functions.  The  sciences,  therefore,  are  not  mutually 
exclusive,  and  their  boundaries  overlap  in  every  direction.  The 
difficulty  in  deciding  what  are  the  most  convenient  boundaries  is 
as  great  with  chemistry  as  with  the  other  groups.  In  particular, 
the  line  which  divides  physics  and  chemistry  from  one  another  is 
often  difficult  to  draw.  It  is  assumed,  however,  that  the  reader 
is  already  familiar  with  the  elements  of  physics,  and  so,  in  place 
of  entering  upon  an  academic  discussion  of  the  nature  of  this  line, 
we  shall  allow  its  location  to  emerge  as  we  proceed. 

The  same  principle  of  grouping  is  pursued  within  each  field. 
Thus  the  preparation  and  properties  of  chemical  compounds  is  called 
descriptive  chemistry.  Again,  the  means  that  have  been  devised 
for  recognizing  the  components  of  mixtures  or  compounds  and 
measuring  their  quantities  constitute  the  several  branches  of  analy- 
sis. The  subdivisions  of  chemistry  of  this  kind  are  numerous. 

The  ideal  in  view  in  thus  classifying  the  content  of  a  science  is  to 
convert  it  into  an  organized  body  of  knowledge.  The  various  ways 
used  to  organize  the  facts  of  a  science  will  be  presented  in  detail  as 
opportunity  offers.  These  ways  constitute  what  is  called  the  scien- 
tific method.  2  ::,::;  :  .-•>.  :••  ;v«  ,\  :  ::** 


2  COLLEGE   CHEMISTRY 

Chemical  Phenomena:  Their  Most  Obvious  Character- 
istic.—  Chemistry  being  primarily  an  experimental  science,  we 
can  best  make  our  first  approach  to  it  through  the  consideration  of 
some  familiar  chemical  phenomena  which  shall  serve  as  illustrations. 

When  clean  iron  is  exposed  to  air  and  moisture  it  becomes  rusty, 
first  on  the  surface,  and  finally  throughout  its  whole  mass.  The  iron 
with  which  we  start  is  a  dark-gray,  metallic-looking  substance,  which 
has  a  high  specific  gravity  (about  7.5),  is  ductile,  has  great  tenacity, 
and  is  readily  attracted  by  a  magnet.  Rust,  on  the  other  hand,  is  a 
reddish  or  brownish  substance  which  has  an  earthy  appearance,  is 
much  lighter  specifically  than  iron  (sp.  gr.  about  4.5),  is  brittle,  and 
is  not  attracted  by  a  magnet.  The  phenomenon  seems  to  have  con- 
sisted in  the  gradual  substitution  of  the  second  of  these  substances 
for  the  first. 

The  chemical  fact  here  is  that  iron,  when  in  contact  with  air  and 
moisture,  gives  rust.  Consideration  will  show,  however,  that  we  do 
not  perceive  this  fact  except  by  noting  the  physical  qualities  of  the 
original  and  of  the  final  materials.  The  data  describing  a  chemical 
phenomenon  consist  in  an  enumeration  of  physical  observations, 
like  the  above  physical  properties  of  iron  and  rust.  It  is,  indeed,  an 
invariable  characteristic  of  every  chemical  phenomenon  that  our 
information  is  all  derived  from  physical  study  of  the  materials. 
Thus,  when  a  candle  burns,  it  undergoes  a  chemical  change,  and  we 
observe  a  solid,  waxy  substance  disappearing  progressively  from 
view.  A  closer  study  of  the  facts  shows  that  a  mixture  of  gases  is 
rising  from  the  flame,  and  we  find  that  the  material  of  the  candle  is 
contained  in  this.  Here,  again,  we  use  physical  means  of  observa- 
tion. 

It  is  further  true  of  these,  as  of  all  chemical  phenomena,  that  the 
physical  properties  of  the  original  and  those  of  the  final  substances 
are  invariably  entirely  different.  We  observe  also  that  this  sort  of 
transformation  is  always  abrupt.  As  the  change  spreads,  the  minute 
fragment  of  iron  of  one  moment  becomes  the  minute  fragment  of 
rust  of  the  next,  and  is  endowed  at  once  with  all  the  new  properties 
in  full  measure.  No  part  of  the  mass  loses  its  magnetic  qualities 
completely,  for  example,  before  it  has  reached  the  limit  of  change  in 
color  or  tenacity.  Since  we  are  compelled  to  define  the  species  of 
matter  by  their  constant  physical  properties  (see  pp.  20,  23),  the 
possession  by  twaor  mqrebodie^of  permanently  different  properties 


INTRODUCTORY  I  3 

constitutes  them  ipso  facto  samples  of  different  materials.  We  may, 
therefore,  say,  more  briefly,  that  the  products  of  a  chemical  change 
are  recognized  at  once  to  be  new  and  different  substances. 

Thus  the  most  obvious  characteristic  of  a  chemical  phenomenon  is: 
that  all  the  physical  properties  of  the  substances  alter;  that  this  alter- 
ation is  abrupt,  so  that,  in  fact,  the  products  are  different  substances;  and 
that  the  recognition  and  study  of  such  a  phenomenon  is  accomplished 
entirely  by  observations  of  a  physical  nature.  Other  characteristics, 
less  obvious,  but  equally  constant,  will  presently  be  brought  to  light. 

The  chemical  phenomena  which  are  familiar  are  innumerable. 
The  burning  of  illuminating-gas,  the  coagulation  of  the  white  of  an 
egg,  the  decay  of  animal  and  vegetable  matter,  the  action  of  hard 
water  on  soap,  and  the  slaking  of  quicklime,  are  examples.  The 
reader  should  analyze  these  changes,  applying  the  above  criteria, 
and  see  for  himself  why  they  are  classed  as  chemical. 

So  far  as  this  first  characteristic  is  concerned,  some  merely  physi- 
cal phenomena  will  be  recalled,  which  are  not  unlike  the  rusting  of 
iron.  Thus,  when  water  is  raised  in  temperature,  its  properties  begin 
to  alter:  it  becomes  more  mobile,  its  specific  gravity  changes,  its 
refracting  power  for  light  is  modified,  and  so  forth.  At  first,  how- 
ever, these  changes  proceed  by  imperceptible  gradations  and  lack  the 
abruptness  of  chemical  change.  But  when  100°  C.  is  reached  the 
water  passes  into  steam,  and  the  whole  of  the  properties  of  this  gas 
are  different  from  those  of  water.  Conversely,  by  cooling,  the  water 
may  be  converted  into  ice,  and  again  there  is  an  abrupt  and  profound 
change  in  properties.  On  this  account,  some  chemists  classify 
"  changes  of  state  of  aggregation  "  as  chemical  phenomena.  More 
usually,  however,  ice,  water,  and  steam  are  regarded  as  identical, 
and  not  different  substances. 

Fact  and  Generalization  or  Law.  —  If  the  preceding  section 
be  reexamined,  it  will  be  seen  that  a  number  of  facts  are  mentioned 
in  it.  We  begin  the  organization  of  knowledge  according  to  the 
scientific  method  by  trying  to  determine  the  facts.  Thus,  we  find 
some  specimens  of  iron  are  variously  colored,  and  some  are  brittle. 
Examination  shows,  however,  that  the  former  peculiarities  are  due 
to  paint,  for  example,  and  the  latter  to  the  presence  of  carbon  and 
other  foreign  materials  in  the  iron  (cast  iron).  Finally,  we  ascer- 
tain the  facts  that  iron  itself  is  gray  and  tough. 


4  COLLEGE   CHEMISTRY 

Putting  together  a  statement  like  that  appearing  in  heavy  type 
above,  is  the  second  step.  We  examine  facts  of  a  like  kind,  or  per- 
taining to  like  phenomena,  to  see  whether  any  general  statement 
can  be  made  that  will  cover  some  feature  common  to  the  whole  of 
them.  In  the  above  illustrations,  after  settling  the  intrinsic  prop- 
erties of  several  substances,  and  then  determining  the  facts  about 
the  way  in  which  these  properties  are  affected  under  certain  circum- 
stances, we  decide  that,  when  all  the  properties  change  abruptly  in 
a  permanent  way,  the  cases  in  which  this  takes  place  shall  consti- 
tute a  distinct  class,  and  we  call  them  cases  of  chemical  change. 
The  statement  which,  in  a  few  words  or  phrases,  sums  up  those 
features  of  all  the  phenomena  of  like  kind  which  are  constant,  is 
called  a  generalization.  Often  it  is  called  a  law,  and  sometimes  a 
principle.  A  generalization  or  law  in  chemistry,  therefore,  is  a  brief 
statement  describing  some  constant  mode  of  behavior.  The  most  im- 
portant laws,  like  that  describing  the  behavior  of  gases  when  com- 
pressed (Boyle's  law),  are  usually  connected  with  the  names  of  the 
men  by  whom  the  relationships  were  discovered  and  the  general- 
izations formulated. 

The  Explanation  of  Rusting.  —  In  considering  the  complete 
disguise  which  is  assumed  by  a  substance  that  has  undergone  a 
chemical  change,  like  rusting,  the  first  question  which  arises  is:  Can 
we  in  any  way  account  for  this  great  and  permanent  change  in 
properties  which  very  generally  distinguishes  the  chemical  pheno- 
menon from  the  physical?  Some  additional  facts  will  be  required 
before  we  can  answer  this  question.  It  has  long  been  known  that 
many  metals  besides  iron  undergo  an  alteration  similar  to  rusting, 
and  that,  where  the  transformation  does  not  occur  spontaneously, 
heating  prompts  it.  The  first  fact  which  seemed  to  throw  light  on 
the  subject  was  the  observation  that  a  piece  of  metal  becomes 
heavier  when  it  rusts.  This  was  noticed  as  early  as  1630,  by  a 
French  physician,  Jean  Rey.  His  work  was  done  chiefly  with  lead 
and  tin,  the  former  of  which  gives  a  dirty  yellow,  and  the  latter,  a 
white  powder,  on  rusting.  He  inferred  correctly  from  his  experi- 
ments that  contact  with  the  air  had  something  to  do  with  this 
chemical  change.  Other  investigations  on  the  same  subject  were 
made  by  Boyle  (1627-1691),  Mayow  (1645-1679),  and  Hooke 
(1635-1703),  in  England,  and  they  led  to  the  same  conclusion.  The 


INTRODUCTORY  I  5 

increase  in  weight  was  to  be  accounted  for,  therefore,  by  the  suppo- 
sition that  some  material  from  the  air  had  been  added  to  the  metal 
during  the  process.  In  other  words,  iron,  for  example,  was  one 
substance  composed  of  iron  only,  and  rust  was  another  substance 
composed  of  iron  and  some  other  material  taken  from  the  atmos- 
phere; and,  the  two  substances  being  different  in  composition,  their 
properties  might  naturally  be  expected  to  differ.  The  substance 
taken  from  the  air  was  subsequently  named  oxygen. 

The  fact  that  foreign  matter  was  actually  gained  by  a  body  during 
the  process  of  rusting  was  not  generally  accepted,  however,  until  it 
was  demonstrated  anew  by  Lavoisier  (1774)  a  century  and  a  half 
later.  He  showed  that  a  portion  of  the  air  really  disappeared  when 
tin  rusted,  and  that  the  increase  in  weight  of  the  tin  corresponded 
with  the  loss  in  matter  of  the  air. 

Explanation  in  Science.  —  The  word  explanation,  which  was 
employed  repeatedly  in  the  last  section,  is  used  in  science  as  the 
name  of  a  definite  process.  It  stands  simply  for  a  description  in 
greater  detail.  Thus,  when,  to  the  acquaintance  with  the  outward 
manifestations  of  rusting,  we  are  able  to  add  the  further  description 
that  it  is  produced  by  the  union  of  oxygen  from  the  air  with  iron,  we 
feel  increased  satisfaction,  and  we  say  that  an  explanation  has  been 
found. 

There  is  no  such  thing  as  a  final  explanation,  however.  At  the 
very  next  step,  when  we  ask  why  the  union  of  oxygen  and  iron  pro- 
duces a  body  that  is  red  and  non-magnetic,  we  are  compelled  to  say 
we  do  not  know. 

An  explanation  in  science  never  professes  for  a  moment  to  give 
the  reasons  for  any  occurrence.  We  simply  do  not  know  why  be- 
havior in  nature  is  as  it  is. 

TJiree  Illustrative  Chemical  Phenomena.  —  Since  oxygen  is 
an  invisible  gas,  the  demonstration  that  rusting  consists  in  the  union 
of  this  gas  with  a  metal,  requires  somewhat  complicated  apparatus. 
The  next  illustration,  while  lacking  historical  interest,  is  simpler, 
because  both  substances  are  visible,  and  are  easily  handled. 

Iron  unites,  not  only  with  oxygen  from  the  air,  but  also,  with 
almost  equal  ease,  with  sulphur.  To  study  the  behavior  of  the 
materials  we  must  know  their  properties.  The  properties  of  iron 


6  COLLEGE   CHEMISTRY 

we  have  already  enumerated  (p.  2).  Sulphur  is  a  pale-yellow  sub- 
stance of  low  specific  gravity  (sp.  gr.  2).  It  is  easily  melted  (rn.-p. 
114.5°C.),  and,  although  it  does  not  dissolve  in  water,  it  may  be 
dissolved  completely  in  carbon  disulphide.  It  crystallizes  in  rhom- 
bic forms  (Fig.  1).  Now,  when  iron  filings 
and  powdered  sulphur  are  rubbed  together 
in  a  mortar,  the  product,  although  it  has  a 
different  color  from  either  of  the  constiu- 
ents,  is  still  really  composed  of  the  two  origi- 
nal kinds  of  matter,  side  by  side.  With  the 
help  of  a  microscope  and  a  needle,  they  can 
be  picked  apart  completely.  By  manipulation 
FIG.  i  of  the  mixture  with  a  magnet,  we  may  remove 

some  of  the  iron  without  much  difficulty. 

Again,  using  the  solubility  of  sulphur  in  carbon  disulphide  and 
the  insolubility  of  iron,  we  may  shake  a  portion  of  the  mixture 
with  this  liquid  in  a  test-tube,  and  so  dissolve  out  the  sulphur 
(Fig.  2).  When  the  contents  of  the  tube  are  poured  on  to  a  filter 
(Fig.  3),  the  liquid  (the  filtrate)  runs 
through,  carrying  all  dissolved  matter 
with  it,  and  leaving  undissolved  mat- 
ter behind.  The  filtrate  may  be 
allowed  to  evaporate  (Fig.  4),  when 
yellow  crystals  of  rhombic  outline 
will  be  found  to  be  the  sole  residue. 
The  dark  material  remaining  on  the 
filter,  when  dry,  is  wholly  attracted 
by  a  magnet.  All  these  facts  con- 
vince us  that  the  properties  of  the 
components  are  still  unaltered,  and 
that,  therefore,  no  chemical  change  FlG  2 

has  occurred. 

If  we  now  put  some  of  the  original  mixture  into  a  test-tube  and 
warm  it,  we  soon  notice  a  rather  violent  development  of  heat  taking 
place,  the  contents  begin  to  glow,  and  what  appears  to  be  a  form  of 
combustion  spreads  through  the  mass.  The  heating  employed  at 
the  start  falls  far  short  of  accounting  for  the  much  greater  heat  pro- 
duced. When  these  phenomena  have  ceased,  and  the  test-tube  has 
been  allowed  to  cool,  we  find  that  it  now  contains  a  somewhat 


INTRODUCTORY  I 


porous-looking,  black  solid.  This  material  is  brittle;  it  is  not  mag- 
netic; it  does  not  dissolve  in  carbon  disulphide;  and  close  exami- 
nation, even  under  a  microscope,  does  not  reveal  the  presence  of 
different  kinds  of  matter.  This  substance  is  known  to  chemists  as 
ferrous  sulphide,  and,  _ 

as  we  see,  its  properties 
are  entirely  different 
from  those  of  the  con- 
stituents. 

A  substance  formed 
in  this  way  by  the 
union  of  other  materials 
is  called  a  compound. 
The  rusts  given  by  vari- 
ous metals  are  there- 
fore compounds  also. 

The  second  illustration 
is  selected  on  account  — 
of  its  historical  interest. 
One  of  the  earliest  chem- 
ical changes  in  which  a 
gas,  recognized  to  be  distinct  from  air,  was  observed  among  the 
products,  was  noticed  by  Priestley  (1774).  The  observation  was 
made  with  mercuric  oxide,  a  bright  red,  rather  ^ep^y  powder. 
When  this  substance  is  heated  (Fig.  5),  we  find  that  a  gas  is  given 
off,  which  is  easily  shown  to  be  different  from  air,  since  a  glowing 
splinter  of  wood  is  instantly  relighted  on  being  immersed  in  it. 
The  gas  is  pure  oxygen.*  We  notice  also  during  the  heating  that  a 

sort  of  mirror  appears  on  the 
sides  of   the  tube.     As  this 
shining    substance    accumu- 
lates it  takes  the  form  of  glob- 
ules, which  may  be  scraped 
together.     It  is,  in  fact,  the 
metal  mercury,  or  quicksilver. 
If  the  heating  continues  long  enough,  the  whole  of  the  red  powder 
eventually  disappears,  and  is  converted  into  these  products. 
The  extensive  nature  of  the  change  in  properties  in  this  case  is 
*  Air  (q.v.)  is  a  mixture  of  which  only  one-fifth  is  oxygen. 


FIG.  3 


FIG.  4 


8 


COLLEGE   CHEMISTEY 


PIG.  5 


evident.     It  should  also  be  observed  that  continuous  heating  is 
required  to  maintain  this  change  in  operation.     It  differs  markedly 

from  the  iron  and  sulphur  case  in  this 
respect.  When  the  flame  is  removed, 
the  evolution  of  oxygen  ceases.  The 
significance  of  this  will  appear  shortly. 
The  third  and  last  example  is  taken 
purposely  in  order  to  illustrate  the 
variety  of  ways  in  which  chemical 
change  may  be  carried  out.  It  is  the 
interaction  of  silver  nitrate  and  sodium 
chloride  (common  salt).  The  substances 
may  be  recognized  by  the  form  of  the 
crystals  of  which  they  consist.  The 
latter  is  composed  of  small  cubes  (Fig. 
6),  while  the  former  presents  a  less  familiar  form  geometrically 
(Fig.  7).  Both  substances  are  capable  of  being  dissolved  in 
water  and,  for  this  experiment,  portions  of  each  substance  are 
shaken  in  separate  vessels  with 
water,  until  none  of  the  solid 
remains.  When  the  solutions  are 
now  poured  together,  we  observe 
that  the  clear  liquids  at  once  be- 
come opaque,  and  that  a  dense 
mass  of  white,  solid  material 
appears  suspended  in  the  mixture  (Fig.  8).  This  white  substance 
consists  of  an  extremely  fine  powder  without 
any  observable  crystalline  form.  We  know 
at  once  that  it  must  represent  a  new  sub- 
stance, since  it  would  not  have  appeared  had 
it  been  soluble  in  water  like  the  two  mate- 
rials from  which  it  was  made.  We  continue 
adding  the  one  liquid  gradually  to  the  other 
until  no  further  formation  of  this  solid  takes 
place,  and  then  stop.  By  filtration  (Fig.  3), 
we  obtain  the  insoluble  material  (the  precipi- 
tate) upon  the  filter  paper,  and  the  clear 
liquid  (the  filtrate)  passes  through  and  is  caught  in  the  vessel  below. 
We  are  confronted  with  two  possibilities:  either  both  the  original 


FIG.  6 


FIG.  7 


FIG.  8 


INTRODUCTORY  I  9 

materials  have  come  together  to  form  one  white  insoluble  material, 
or  some  other  product  (or  products)  may  be  present  in  addition  to 
it.  In  the  latter  case,  search  must  evidently  be  made  in  the  liquid. 
By  evaporating  the  filtrate  in  a  suitable  vessel  (Fig.  4),  we  find  that 
the  second  assumption  represents  the  fact,  for  a  considerable  quan- 
tity of  a  white  crystalline  substance  remains.  The  homogeneous 
character  of  this  shows  that  there  was  but  one  product  in  solution, 
while  the  same  property  of  the  precipitate  shows  that  there  are  but 
two  products  altogether. 

The  insoluble  material  is  composed  of  silver  and  chlorine,  and  it 
is  known  as  silver  chloride.  Like  some  other  compounds  of  silver, 
it  darkens  on  exposure  to  light,  turning  first  purple  and  then  brown, 
and  being  decomposed  by  this  agency  into  its  constituents.  The 
chlorine,  a  gas,  escapes,  and  a  brown  powder  consisting  of  finely 
divided  silver  remains.  The  soluble  solid 
obtained  from  the  filtrate,  we  recognize  as 
identical  with  a  mineral,  sodium  nitrate, 
which  is  found  in  Peru.  Its  crystals  are 
rhombohedral  (Fig.  9).  They  resemble  cubes 
which  have  been  slightly  distorted  by  pressing 
inwards  two  opposite  corners.  This  change 

presents  several  features  which  distinguish  it  from  the  previous 
ones:  it  is  much  more  complex;  it  takes  place  in  the  presence  of 
water;  it  requires  no  heating  for  its  promotion;  and  the  change  is 
complete  the  instant  the  materials  have  been  mixed,  while  the 
others  required  a  good  deal  of  time  for  their  accomplishment. 

The  Kinds  of  Chemical  Change.  —  Having  these  three  cases 
before  us,  —  and  they  are  types  of  most  chemical  phenomena,  —  we 
now  proceed  to  analyze  them,  and  so  take  another  step  towards 
explaining  (i.e.,  describing  in  detail)  the  nature  of  a  chemical  phe- 
nomenon. In  the  iron  and  sulphur  experiment  (p.  6),  two  mate- 
rials were  used,  and  a  different  one  with  new  properties  was 
produced.  Here,  in  chemical  language,  the  first  two  substances 
united  or  underwent  combination:* 

Iron  4-  Sulphur  gave  Ferrous  sulphide. 

The  rusting  of  iron,  lead,  and  tin  belongs  also  to  this  class.  In  the 
second  illustration  (p.  7),  one  material  was  used,  and  it  was  driven 

*  The  word  mixed  is  used  only  of  substances  which  are  mingled,  physically, 
and  not  combined. 


10  COLLEGE  CHEMISTRY 

apart  by  heating  so  that  two  new  ones  arose.  A  chemical  pheno- 
menon of  this  kind  is  called  a  decomposition: 

Mercuric  oxide  gave  Mercury  +  Oxygen. 

The  last  was  the  most  complex.  The  substances,  with  the  constit- 
uents of  each,  were  as  follows: 

Silver  nitrate  +  Sodium  Chloride  gave  Silver  Chloride  +  Sodium  nitrate. 
(Silver,  nitrogen,  oxygen)  (Sodium,  chlorine)  (Silver,  chlorine)  (Sodium,  nitrogen,  oxygen) 

Here  the  constituents  of  both  the  ingredients  became  separated 
(decomposition),  and  the  products  of  decomposition  mated  them- 
selves differently  (combination).  Thus,  for  example,  the  silver  of 
one  ingredient  combined  with  the  chlorine  of  the  other.  A  third 
variety  of  chemical  change  consists,  therefore,  in  the  concurrence  of 
both  of  the  first  two  kinds  in  the  same  action.  As  several  different 
varieties  of  this  sort  of  complex  redistribution  of  material  can  be 
distinguished,  a  distinct  name  is  given  to  each.  The  present  is 
called  a  double  decomposition. 

When  we  encounter  other  chemical  changes,  we  shall  find  the 
extent  of  the  stride  we  have  taken  in  this  critical  analysis  of  a  few 
examples.  It  will  then  appear  that  there  are  only  a  few  changes 
which  cannot  be  placed  in  one  of  these  three  categories,  and  these 
all  belong  to  a  fourth  class.  It  occasionally  happens,  especially  in 
the  case  of  compounds  of  carbon,  that  one  single  kind  of  material 
turns  into  another  single  kind  of  material.  Nothing  is  added  and 
nothing  removed,  yet  the  new  substance  has  different  properties  in 
every  respect  from  the  old.  Most  of  those  substances  whose 
transformation  is  definitely  assigned  to  this  class  contain  several 
constituents.  Now,  we  have  seen  that  when  substances  have  dif- 
ferent properties  they  differ  also  in  their  constituent  materials. 
Carrying  out  this  idea,  the  hypothesis  (or  suggestion)  has  been  made 
that,  since  here  also  the  properties  change,  there  must  be  some  read- 
justment of  the  material,  even  in  cases  like  this.  Hence,  we  desig- 
nate changes  of  this  kind  internal  rearrangements.  The  composition 
of  the  material  is  unaltered,  so  we  suppose  its  constitution  to  have 
become  different.  If  the  chemists  ever  decide  to  regard  the  change 
of  water  into  ice  or  steam  as  a  chemical  phenomenon,  all  changes 
of  state  of  aggregation  would  be  placed  in  this  fourth  class.  It  is 
here  only  that  the  boundary  between  physics  and  chemistry  is  at 
present  difficult  to  define. 


INTRODUCTORY  I  11 

The  Second  Characteristic  of  Chemical  Phenomena.  —  We 

are~now  able  to  make  another  generalization  (or  condensed  state- 
ment of  fact) :  In  chemical  phenomena  substances  enter  into  combina- 
tion, come  out  of  combination,  or  change  their  associates  in  combination 
or  their  state  in  the  compound.  In  other  words,  the  material  changes 
its  composition  or  its  constitution. 

This  generalization  furnishes  an  explanation  of  the  former  (p.  3) 
to  a  certain  extent.  Each  individual  substance  has  its  own  compo- 
sition and  its  own  set  of  physical  properties.  Hence,  when,  by 
chemical  transformation,  substances  of  new  and  different  composi- 
tion are  produced,  the  materials  must  simultaneously  acquire  the 
specific  properties  of  the  new  substances. 

A  chemical  phenomenon  is  named  an  action  or  interaction.  It  is 
also  often  called  a  chemical  reaction.  Thus,  we  say  that  four 
classes  of  reactions  have  been  described  above. 

Summary.  —  Thus  far,  we  have  learned  that  chemistry  deals 
with  the  changes  in  composition  and  constitution  which  substances 
undergo,  and  with  the  alteration  in  properties  which  accompanies 
and  gives  evidence  of  those  changes. 

Exercises.  —  1.  Take  one  by  one  the  words  or  phrases  printed 
in  black  type  and  the  titles  of  the  sections  in  this  chapter,  and  en- 
deavor to  recollect  what  you  have  read  about  each.  In  each  case 
try,  (a)  to  recall  the  meaning  and  to  state  it  in  your  own  words; 
(6)  to  recall  the  facts  associated  with,  and  the  reasoning  which  led 
up  to  the  point  in  question;  (c)  to  recall  examples  illustrating  the 
conception  and  to  apply  the  conception  in  detail  to  each  example. 
Whenever  memory  fails  to  give  a  perfectly  clear  report  of  the  mat- 
ter in  hand,  the  text  must  be  read  and  reread  until  the  essential 
point  can  be  repeated  from  memory. 

Use  the  same  method  in  all  future  chapters.  A  useful  practice  is 
to  employ  a  pencil  as  you  read  and  to  underline  systematically  all 
the  important  facts  and  statements,  and  then  to  go  back  and  apply 
to  each  marked  place  the  process  described  above. 

2.  Define  the  following  and  state  how  each  is  measured:  specific 
gravity,  ductility,  tenacity,  brittleness. 

3.  Take  the  chemical  phenomena  mentioned  on  p.  3,  par.  2,  and 
enumerate  the  physical  properties  of  the  substances  before  and 
after  the  chemical  change. 

4.  Define  and  illustrate:  filtrate,  mixture,  compound,  precipitate. 


CHAPTER  II 
INTRODUCTORY   H 

IF  we  now  return  to  the  three  illustrations  of  chemical  phenomena 
which  we  have  been  studying  (pp.  5-9),  we  shall  find  a  new 
question  arising  naturally  out  of  them.  This  is,  whether  the 
mass  of  the  materials  is  altered,  as  are  the  other  attributes,  in 
these  chemical  changes. 

Third  Characteristic  of  Chemical  Phenomena:  Conser- 
vation of  Mass.  —  The  most  painstaking  chemical  work  seems  to 
show  that,  if  all  the  substances  concerned  in  the  chemical  change 
are  weighed  before  and  after  the  change,  there  is  no  evidence  of  any 
alteration  in  the  quantity  of  matter.  The  two  weights,  represent- 
ing the  sums  of  the  constituents  and  of  the  products  respectively, 
are,  indeed,  never  absolutely  identical,  but  the  more  careful  the 
work  and  the  more  delicate  the  instrument  used  in  weighing,  the 
more  nearly  do  the  values  approach  identity.  We  are  able  to 
state,  therefore,  as  a  third  characteristic  of  all  chemical  phenomena, 
that  the  mass  of  a  system  is  not  affected  by  any  chemical  change  within 
the  system. 

This  statement  simply  means  that  the  great  law  of  the  conserva- 
tion of  mass  holds  true  in  chemistry  as  it  does  in  physics.  Chemi- 
cal changes,  thoroughgoing  as  they  are  in  respect  to  all  other  prop- 
erties, do  not  affect  the  mass;  an  element  carries  with  it  its  weight, 
entirely  unchanged,  through  the  most  complicated  chemical  trans- 
formations. This  is  the  only  attribute  which  persists. 

Superficial  observation,  as  of  a  growing  tree,  might  seem  to  give 
evidence  of  the  very  opposite  of  conservation  of  matter.  But  here 
the  carbon  dioxide  gas  in  the  air,  the  most  important  source  of 
nourishment  for  plants,  is  overlooked.  Similarly  the  gradual  dis- 
appearance of  a  candle  by  combustion  seems  to  illustrate  the 
destruction  of  matter.  But  if  we  catch  the  gases  which  rise  through 
the  flame,  we  find  that  the  gases  weigh  even  more  than  the  part  of 

12 


INTRODUCTORY  II  13 

the  candle  which  has  been  sacrificed  in  making  them.  When  we 
take  account  of  the  weight  of  the  oxygen  obtained  from  the  air 
which  sustains  the  combustion,  we  find  that  there  is  really  neither 
loss  nor  gain  in  weight.  If  we  carry  out  chemical  changes  in  closed 
vessels,  which  permit  neither  escape  nor  access  of  material,  we  find 
that  the  weight  does  not  alter. 

Physical  Accompaniments  of  Chemical  Change.  —  Study 
of  the  three  typical  chemical  changes  described  in  the  last  chapter 
may  now  be  resumed,  in  order  to  see  whether  anything  further  of  a 
general  nature  is  characteristic  of  such  phenomena.  We  recall  at 
once  that  a  prominent  feature  of  the  union  of  iron  and  sulphur  was 
the  heat  which,  as  shown  by  the  glow  spreading  through  the  mass, 
seemed  to  be  developed  after  the  action  was  once  started.  It  is 
found  that  many  chemical  changes  are  like  this  one,  in  exhibiting 
simultaneously  the  production  of  very  perceptible  amounts  of  heat. 
On  the  other  hand,  the  decomposition  of  mercuric  oxide,  as  was 
pointed  out  (p.  8),  owed  its  continuance  to  the  persistent  applica- 
tion of  heat,  and  ceased  so  soon  as  the  source  of  heat  was  with- 
drawn. Here,  apparently,  heat  was  consumed  during  the  progress 
of  the  change,  and  the  chemical  action  was  limited  by  the  amount  of 
heat  supplied.  The  production  or  consumption  of  heat  may,  there- 
fore, be  a  feature  of  chemical  change. 

In  the  iron  and  sulphur  case,  as  in  other  chemical  actions  where 
the  heat  developed  is  great,  light  also  was  given  out.  In  the  last  of 
the  three  actions,  on  the  other  hand,  we  obtained  a  substance  (silver 
chloride),  which  may  be  kept  for  any  length  of  time  in  the  dark, 
but,  by  the  action  of  sunlight  is  broken  up  into  its  constituents 
(p.  9).  It  would  appear,  therefore,  that  light  may  be  given  out  or 
used  in  connection  with  chemical  change.  Noting  these  facts  stim- 
ulates us  to  look  for  other  similar  accompaniments  of  change  in 
composition. 

If  we  dip  two  wires  from  a  battery  or  dynamo  into  a  solution  of 
nitrate  of  silver  (Fig.  10),  such  as  was  used  in  the  third  experiment, 
we  observe  the  instant  production  of  a  coating  of  silver  on  the  nega- 
tive wire.  By  preparing  the  solution  properly  and  allowing  the 
electricity  to  flow  through  it  for  a  sufficient  length  of  time,  all  of  the 
compound  can  be  decomposed  and  all  its  silver  deposited.  It  is 
needless  to  say  that  this  release  of  the  silver  from  chemical  combina- 


14 


COLLEGE   CHEMISTRY 


FIG.  10 


tion,  and  liberation  of  the  metal  at  the  electrode,  goes  on  only  so 
long  as  the  current  of  electricity  is  employed,  and  that  electricity  is 

consumed  in  the  pro- 
cess.    Very  many  sub- 
~T5>h"N  //S*  stances  can  be  decom- 

8 O=H*  NV  posed  in  this  way. 

The  inverse  of  this 
is  likewise  familiar.  If 
we  place  in  dilute  sul- 
phuric acid  a  stick  of 
the  metal  zinc,  we  find 
that  a  gas  is  given  off 
rapidly  (Fig.  11),  that 
the  zinc  gradually  dis- 
solves, and  that  a  large 
amount  of  heat  is  de- 
veloped. Under  favor- 
able circumstances,  the 
liquid  may  even  rise  Spontaneously  to  the  boiling-point.  This  form 
of  the  action  produces  heat.  If,  however,  we  attach  the  same  stick 
of  zinc  to  a  copper  wire,  and,  having  pro- 
vided a  plate  of  platinum  also  connected 
with  a  wire,  immerse  the  two  simultaneously 
in  the  acid  (Fig.  12),  then  a  galvanometer, 
with  which  the  wires  are  connected,  shows 
at  once  the  passage  of  a  current  of  electricity 
round  the  circuit.  Exactly  the  same  chemical 
change  goes  on  as  before.  The  sole  difference 
is  that  the  gas  appears  to  arise  from  the 
surface  of  the  platinum.  It  is  easy  to  show, 
however,  that  the  platinum  by  itself  is  not 
acted  upon  by  dilute  acids,  and,  in  this  case, 
undergoes  no  change  whatever;  it  serves 
simply  as  a  suitable  conductor  for  the  elec- 
tricity. Here,  then,  in  place  of  the  heat 
which  the  first  plan  produced,  we  get  elec- 
tricity. The  arrangement  is,  in  fact,  a 
battery-cell,  for  a  battery  is  a  system  in  which  a  chemical  action 
which  would  otherwise  give  heat  furnishes  electricity  instead.  Thus, 


FIG.  ll 


INTRODUCTORY  II 


15 


electricity  may  be  consumed  or  produced  in  connection  with  a  change 
in  composition. 

Even  violent  rubbing  in  a  mortar,  in  the  case  of  some  substances, 
can  effect  an  appreciable  amount  of  decomposition  in  a  few  minutes. 
In  this  way  silver  chloride  can  be  separated  into  silver  and  chlorine, 
just  as  by  light.  It  is  the  mechanical  energy  which  is  the  agent,  and 
part  of  it  is  consumed  in  producing  the  change,  and  only  the  balance 


FIG.  12 


appears  as  heat.  Conversely,  the  production  of  mechanical  energy,  as 
the  result  of  chemical  change,  is  seen  in  the  behavior  of  explosives 
and  in  the  working  of  our  muscles.  Thus,  mechanical  energy  may  be 
used  up  or  produced  in  chemical  changes. 

Summing  up  our  experience,  we  may  state  that  no  change  in  com- 
position occurs  without  some  accompaniments,  such  as  the  produc- 
tion or  consumption  of  heat,  light,  electricity,  or,  in  some  cases, 
mechanical  energy. 

Classification  of  the  Accompaniments  of  Change  in  Com- 
position: Energy.  —  The  problem  of  classifying  (i.e.,  placing  in  a 
suitable  category)  things  like  heat,  light,  and  electricity  has  occu- 
pied much  attention.  They  do  not  possess  mass.  In  all  changes  in 
composition,  one  of  these  natural  accompaniments  is  given  out  or 
absorbed,  sometimes  in  great  amount,  yet,  in  none  is  any  alteration 
in  weight  observed.  There  are  many  things  which  are  real,  how- 


16 


COLLEGE   CHEMISTRY 


ever,  even  if  they  are  not  affected  by  gravitation.     In  the  present 
instance  we  reason  as  follows: 

A  brick  in  motion  is  different  from  a  brick  at  rest.  The  former 
can  do  some  things  that  the  latter  cannot.  Furthermore,  we  can 
easily  make  a  distinction  in  our  minds.  The  brick  can  be  deprived  of 
the  motion  and  be  endowed  with  it  again.  Thus,  we  can  get  the  idea 
of  motion  as  a  separate  conception.  Similarly,  we  observe  that  a 
piece  of  iron  behaves  differently  when  hot,  and  when  cold,  when 
bearing  a  current  of  electricity,  and  when  bearing  none.  We  con- 
ceive then  of  the  brick  or  the  iron  as  having  a  certain  amount  and 
kind  of  matter  which  is  unalterable,  and  as  having  motion,  heat,  or 
electricity  added  to  this  or  removed.  Thus,  we  describe  our  obser- 
vations by  using  two  categories,  one  of  which  includes  the  various 
kinds  of  matter,  and  the  other,  various  things  whose  association  with 
matter  seems  to  be  invariable  and  is  often  so  conspicuous. 

At  first  sight,  these  accompaniments  of  matter  seem  to  be  quite 
unrelated.  But  a  relation  between  them  can  be  found.  If  the  heat 

of  a  Bunsen  flame  or  of  the  sun  is 
brought  under  a  hot-air  motor  (Fig.  13) 
violent  motion  results.  Again,  if  the 
motor  is  connected  with  a  dynamo, 
electricity  may  be  generated.  Still 
again,  if  the  current  flows  through  an 
incandescent  lamp,  heat  and  light  are 
evolved.  Conversely,  when  motion  is 
impeded  by  a  brake,  heat  appears. 
When  a  current  of  electricity  is  run 
through  the  dynamo,  motion  results. 
But  the  most  significant  facts  are  still 
to  be  mentioned.  The  heat  absorbed 
by  the  motor  is  found  to  be  greater 
when  the  machine  is  permitted  to 
move  and  do  work,  than  when  it  is 
not.  Thus,  it  is  found  that  when  work  is  done  some  heat  dis- 
appears, and  this  heat  is,  in  fact,  transformed  into  work.  Similarly, 
when  the  poles  of  the  dynamo  are  properly  connected  and  electricity 
is  being  produced,  and  only  then,  motion  is  used  up.  This  is  shown 
by  the  effort  required  to  turn  the  armature  under  these  circum- 
stances, and  the  ease  with  which  it  is  turned  when  the  circuit  is 


FIG.  13 


INTRODUCTORY   II  17 

open.  So,  with  a  conductor  like  the  filament  in  the  lamp,  unless  it 
offers  resistance  to  the  current  and  destroys  a  sufficient  amount 
of  electricity,  it  gives  out  neither  light  nor  heat.  Finally,  motion 
gives  no  heat  unless  the  brake  is  set,  and  effort  is  then  demanded 
to  maintain  the  motion.  These  experiences  lead  us  to  believe  that 
we  have  here  a  set  of  things  which  are  fundamentally  of  the  same 
kind,  for  each  form  can  be  made  from  any  of  the  others.  We  have, 
therefore,  invented  the  conception  of  a  single  thing,  of  which  heat, 
light,  electricity,  and  motion  are  forms,  and  to  jt  we  give  the  name 
energy :  energy  is  work  and  every  other  thing  which  can  arise  from 
work  and  be  converted  into  work  (Ostwald). 

Closer  study  shows  that  equal  amounts  of  electrical  or  mechanical 
energy  always  produce  equal  amounts  of  heat.  There  is  never  ob- 
served any  loss  in  any  of  the  transformations  of  energy  any  more 
than  in  the  transformations  of  matter.  Hence,  J.  R.  Mayer  (1842), 
Colding  (1843),  and  Helmholtz  (1847)  were  led  independently  to  the 
conclusion  that  in  a  limited  system  no  gain  or  loss  of  energy  is  ever 
observed.  This  brief  statement  of  the  results  of  many  experiments 
is  called  the  law  of  the  conservation  of  energy. 

Application  of  the  Conception  of  Energy  in  Chemistry.  —  At 

first  sight  it  looks  as  if  the  statement  that  energy  is  conserved  is  not 
applicable  in  chemistry.  Heat  and  electricity,  for  example,  seem  to 
be  produced  and  consumed,  in  connection  with  changes  in  composi- 
tion, in  a  mysterious  manner.  We  trace  light  in  an  incandescent 
lamp  back  to  the  electricity,  and  this  in  turn  to  the  mechanical 
energy,  and  this  again  to  the  heat  in  the  engine.  But  what  form  of 
energy  gave  the  heat  developed  by  the  combustion  of  the  coal  under 
the  boiler,  or  by  the  union  of  iron  and  sulphur  in  our  first  experiment? 
Since  we  do  not  perceive  any  electricity,  light,  heat,  or  motion,  in  the 
original  materials,  and  yet  wish  to  create  an  harmonious  system,  we 
are  bound  to  conceive  of  the  iron  and  the  sulphur,  and  the  coal  and 
the  air,  as  containing  another  form  of  energy,  which  we  call  internal 
energy.  Similarly,  when  heat  is  used  up  in  decomposing  mercuric 
oxide,  or  light  in  decomposing  silver  chloride,  we  regard  the  energy  as 
being  stored  in  the  products  of  decomposition  in  the  form  of  internal 
energy. 

These  conclusions  compel  us,  for  the  sake  of  consistency,  to  think 
of  all  our  materials  as  repositories  of  energy  as  well  as  of  matter, 


18  COLLEGE   CHEMISTRY 

each  of  these  constituents  being  equally  real  and  equally  important. 
A  piece  of  the  substance  known  as  "  iron  "  must  thus  be  held  to 
contain  so  much  iron  matter  and  so  much  internal  energy.  So 
ferrous  sulphide  contains  sulphur  matter,  iron  matter,  and  internal 
energy.  Thus,  by  a  substance  we  mean  a  distinct  species  of  matter, 
simple  or  compound,  with  its  appropriate  proportion  of  internal 
energy. 

In  the  course  of  this  discussion  it  has  become  clear  that  a  fourth 
characteristic  of  chemical  phenomena  is  that,  besides  a  change  in  the 
state  of  the  matter,  there  is  always  an  alteration  in  the  amount  of  internal 
energy  in  the  system.  This  alteration  consists  in  the  production  of  inter- 
nal energy  from,  or  the  transformation  of  internal  energy  into,  an  equiva- 
lent amount  of  some  other  form  of  energy. 

The  absorption  or  liberation  of  energy  accompanying  a  chemical 
transformation  of  matter  is  often,  of  the  two,  the  more  important 
feature.  We  do  not  burn  coal  in  order  to  manufacture  carbon 
dioxide  gas.  We  are  glad  to  get  rid  of  the  material  product  through 
the  chimney.  It  is  the  heat  we  want.  We  do  not  employ  zinc  in 
batteries  with  the  object  of  making  zinc  chloride  or  zinc  sulphate. 
So  we  use  the  electrical  energy,  and  throw  the  material  products 
away  when  we  refill  the  jars.  It  is  the  same  with  burning  illumi- 
nating-gas or  magnesium  powder  when  we  want  light,  and  with 
eating  food,  which  we  do,  chiefly,  to  get  energy  to  sustain  our 
activity.  We  do  not  run  electricity  for  hours  into  a  storage  battery 
in  order  to  make  a  particular  compound  (lead  dioxide,  for  example), 
but  in  order  to  save  and  store  the  energy  for  future  use.  In  industry 
and  life  fully  half  the  total  amount  of  chemical  change  involved  is 
set  in  motion  by  us,  solely  on  account  of  the  energy  changes  it 
involves.  But  the  production  of  energy  in  chemical  change  is  not 
only  thus  of  practical  importance;  it  is  also  of  scientific  interest,  as 
will  be  seen  in  the  next  section. 

Chemical  Activity.  —  Other  things  being  equal,  actions  in  which 
there  is  a  relatively  large  loss  of  internal  energy  proceed  rapidly; 
that  is  to  say,  in  them  a  large  proportion  of  the  material  is  changed 
in  the  unit  of  time.  Those  in  which  less  energy  is  transformed 
proceed  more  slowly.  The  speed  of  the  chemical  change,  and  the 
quantity  of  energy  available  because  of  it,  are  closely  related.  Now, 
we  are  accustomed  to  speak  of  materials  which,  like  iron  and  sulphur, 


INTRODUCTORY   II  19 

interact  rapidly  and  with  liberation  of  much  energy  as  "  chemically 
active."  Thus,  relative  chemical  activity  may  be  estimated,  (1)  by 
observing  the  speed  of  a  change  (see  Speed  of  chemical  actions),  or, 
in  many  cases  (2)  by  measuring  the  heat  developed  (see  Thermo- 
chemistry), or  (3)  by  ascertaining  the  electromotive  force  of  the 
current  the  change  gives,  when  arranged  in  the  form  of  a  battery- 
cell. 

It  is  evident  that  the  chemical  activity  of  a  given  substance  will 
not  be  the  same  towards  all  others.  Thus,  iron  unites  much  more 
vigorously  with  chlorine  than  with  sulphur,  and,  with  identical 
amounts  of  iron,  more  heat  is  liberated  in  the  former  case  than  in  the 
latter.  With  silver,  sodium,  and  many  other  substances,  iron  does 
not  unite  at  all.  One  of  the  tasks  of  the  chemist  is  to  make  such 
comparisons  as  this  (see  Specific  chemical  properties).  Evidently, 
the  substances  containing  the  most  chemical  energy  will  be  in  gen- 
eral the  most  active. 

The  "  Cause  "of  Chemical  Activity.  —  The  reader  will  un- 
doubtedly be  inclined  to  inquire  whether  we  can  assign  any  cause  for 
the  tendency  which  substances  have  to  undergo  chemical  change. 
Why  do  iron  and  sulphur  unite  to  form  ferrous  sulphide,  while  other 
pairs  of  elements  taken  at  random  will  frequently  be  found  to  have 
no  effect  upon  one  another  under  any  circumstances?  The  answer 
is  that  we  do  not  know.  Questions  like  this  have  to  go  without 
answer  in  all  sciences.  What  is  the  cause  of  gravitation?  We  know 
the  facts  which  are  associated  with  the-  word  —  the  fact  that  bodies 
fall  towards  the  earth,  for  example  —  but  why  they  fall  we  are 
unable  to  say.  So,  with  chemical  change,  we  can  state  all  the  facts 
we  know  about  it,  but  even  then  we  cannot  say  why  it  takes  place. 

The  word  cause  was  employed  in  the  heading  of  this  section,  and 
it  will  be  observed  that  no  cause  was  found.  This  is  the  invariable 
rule  in  physical  or  chemical  phenomena.  We  know  of  no  causes,  in 
the  sense  in  which  the  word  is  commonly  employed. 

The  word  cause  has  only  one  definite  use  in  science.  When  we 
find  that  thorough  incorporation  of  the  three  materials  is  needed  to 
secure  good  gunpowder,  we  say  that  the  intimate  mixing  is  a  cause 
of  its  being  highly  explosive.  By  this  we  simply  mean  that  intimate 
mixture  is  a  necessary  antecedent  of  the  result.  A  cause  is  a  condition 
or  occurrence  which  always  precedes  another  condition  or  occurrence. 


20  COLLEGE    CHEMISTRY 

Of  Elements  and  of  Simple  and  Compound  Substances.  — 

If  we  place  before  a  physicist  samples  of  iron,  ferrous  sulphide,  and 
sulphur,  he  will  report  that  there  are  three  absolutely  distinct  sub- 
stances represented,  because  they  show  three  different  sets  of  physical 
properties.  A  chemist,  on  the  other  hand,  while  admitting  the 
accuracy  of  the  report,  in  view  of  the  criterion  used  by  the  physicist, 
which  indeed  he  uses  himself  (c/.  pp.  2-3),  will  insist  on  adding  that 
there  are  only  two  perfectly  distinct  kinds  of  matter  in  the  set, 
because  he  can  make  the  second  from  matter  furnished  by  the  other 
two.  In  a  sense,  chemistry  reduces  the  kinds  of  different  matter  to 
a  much  smaller  number  than  does  physics  or  any  of  the  other  sciences, 
and  so  it  is  the  final  authority  in  all  questions  involving  matter.  By 
the  chemist,  dozens  of  physically  distinct  substances  are  regarded  as 
closely  related  because  they  all  can  be  made  with  iron,  or  when 
decomposed  give  it;  hundreds  are  alike  in  that  sulphur  enters  into 
their  composition;  thousands  are  compounds  of  oxygen.  In  fact, 
the  number  of  kinds  of  matter  which  are  perfectly  distinct  in  the 
strictly  chemical  point  of  view  is  quite  limited. 

The  first  to  put  our  modern  view  into  definite  language  was 
Lavoisier  in  his  Traite  de  Chimie  (1789).  His  work  showed  that 
decomposition  had  its  limits.  Mercuric  oxide  could  be  decomposed 
into  mercury  and  oxygen,  but  no  means  was  found  of  breaking  these 
up  in  turn  and  producing  any  fresh  substances  from  them.  The 
kinds  of  matter  composing  these  simple  materials  he  named  elements. 
The  element  is  to  be  regarded  as  an  ultimate  chemical  individual 
just  as  the  substance  is  the  physical  individual.  The  definition  of 
an  element  is  therefore:  a  distinct  species  of  matter  which  has  not 
been  shown  to  be  composite. 

The  caution  which  prompted  Lavoisier  to  use,  as  he  did,  the  words 
"  has  not  yet  been,"  was  justified  by  the  fact  that  several  substances, 
in  his  time  regarded  as  elementary,  were  afterwards  shown  to  be 
compound.  Thus,  quicklime  was  a  simple  substance  until  Davy, 
in  1808,  prepared  the  metal  calcium  and  showed  that  quicklime  was 
a  compound  of  this  metal  with  oxygen.  Discoveries  similar  to  this 
have  been  made  on  more  than  one  occasion  since. 

Until  recently,  a  body  made  up  of  one  or  more  specific  elements 
had  never  been  found  to  yield  any  simple  substance  different  from 
those  used  in  preparing  it.  In  other  words,  one  element  had  never 
been  turned  into  another.  The  production  of  helium,  neon,  and 


INTRODUCTORY  II  21 

argon  by  the  spontaneous  decomposition  of  the  radium  (q.v.)  emana- 
tion, and  the  formation  of  traces  of  lithium  from  copper  sulphate, 
both  observed  recently  (1907)  by  Sir  William  Ramsay,  may  lead, 
however,  to  a  revision  of  the  conception  of  an  element. 

We  have  seen  (p.  18)  that  all  substances,  that  is,  physical  indi- 
viduals, must  be  thought  of  as  containing  matter  and  energy.  We 
have  now  learned  that  there  are  two  kinds  of  substances,  namely, 
compound  substances  which  contain  two  or  more  elements  together  with 
the  appropriate  amount  of  energy,  and  simple  substances  which  contain 
only  one  element  together  with  a  certain  quantity  of  energy.  Among 
the  substances  which  we  have  been  handling,  iron,  sulphur,  mercury, 
oxygen,  and  hydrogen  are  simple  substances.  On  the  other  hand, 
the  substances  which  we  have  shown  to  be  composite  are  ferrous 
sulphide,  rust,  mercuric  oxide,  silver  nitrate,  and  common  salt.  It 
will  be  seen  that  by  combination  of  a  limited  number  of  elements, 
two,  three,  or  four  together,  in  varying  proportions,  all  the  known 
distinct  substances  might  easily  be  accounted  for.  The  list  of  ele- 
ments whose  individuality  has  been  established  appears  upon  the 
inside  of  the  cover,  at  the  end  of  this  book;  of  these  the  larger  number 
are  not  frequently  encountered.  More  than  99  per  cent  of  terrestrial 
material  is  made  up  of  eighteen  or  twenty  elements,  of  which  the 
quantities  of  the  first  eleven,  as  estimated  by  F.  W.  Clarke,  are  given 
in  the  following  table: 

Oxygen     .    .   .    .49.98  Calcium 3.51  Hydrogen  ....  0.94 

Silicon 25.30  Magnesium     .    .    .2.50  Titanium    .    .    .    .0.30 

Aluminium  ...     7.26  Sodium 2.28  Carbon    ....    .0.21 

Iron 5.08  Potassium   .    .    .    .2.23  99.61 

The  evidence  of  the  spectroscope  shows  that  the  sun  and  stars 
contain  many  of  the  very  same  elements  as  does  the  earth. 

Some  of  the  Fundamental  Ideas  used  by  Chemists  and  the 
Corresponding  Terms.  —  To  the  chemist  it  is  above  all  im- 
portant that  he  should  be  able  to  set  forth  in  unambiguous  terms 
the  phenomena  he  observes  and  the  inferences  he  draws.  To  do 
this  he  requires  some  fundamental  ideas  described  by  suitable 
words.  Many  of  these  ideas  and  terms  we  have  been  employing  in 
order  to  accustom  the  reader  to  their  use.  It' is  now  advisable  to 
take  them  up  more  systematically. 


22  COLLEGE    CHEMISTRY 

Any  particular  specimen  of  matter,  such  as  a  piece  of  sulphur,  a 
portion  of  water,  a  piece  of  ferrous  sulphide,  a  fragment  of  granite,  or 
some  nitrate  of  silver  solution,  we  call  a  body.  There  are  thus  as 
many  bodies  as  there  are  discrete  portions  of  matter.  A  body  may 
be  heterogeneous,  or  made  up  of  visibly  unlike  parts,  as  granite 
and  a  mixture  of  iron  powder  and  sulphur  are;  or  it  may  be  homoge- 
neous, or  alike  in  all  parts,  as  are  pieces  of  sulphur  and  ferrous  sul- 
phide and  portions  of  water  and  silver  nitrate  solution. 

Examination  of  these  homogeneous  bodies  shows  that  the  sulphur, 
ferrous  sulphide,  and  water  differ  from  the  nitrate  of  silver  solution 
in  having  but  one  physical  component,  while  the  last  contains  two 
components,  nitrate  of  silver  and  water,  separable  by  physical 
means.  The  last  is  like  a  mixture,  only  it  is  homogeneous.  Again, 
the  sulphur,  ferrous  sulphide,  and  water  differ  amongst  themselves, 
the  first  being  a  simple  body,  with  one  chemical  constituent,  and  the 
two  others  being  compounds  having  each  two  chemical  constituents. 
We  speak  of  the  components  of  a  mixture"  or  a  solution  because  the 
parts  are  laid  together  and  retain  in  the  former  case  all,  and  in  the 
latter  much,  of  their  identity.  But  of  the  materials  in  a  chemical 
compound  we  use  the  word  constituents  because  the  parts  are  built 
into  each  other  and  have  lost  their  identity.  This  distinction  is  very 
important.  Unless  it  is  understood,  we  cannot  ourselves  clearly 
describe  a  given  body,  or  understand  the  description  when  given 
by  someone  else.  When  we  state  the  result  of  our  study  of  a  body, 
as,  for  example,  some  liquid  or  solid  in  a  test  tube,  we  always  give, 
first,  the  physical  components.  Thus  we  might  say  the  material  was 
composed  of  common  salt  in  solution  in  water,  or  of  a  mixture  of 
sand  and  sugar.  Usually  this  is  all  the  information  that  is  needed, 
since  the  nature  of  the  components  named  is  presumably  known. 
If,  however,  more  information  is  demanded,  then  we  proceed,  next, 
to  name  the  chemical  constituents  of  each  of  the  components.  Thus, 
we  might  add  that  sodium  and  chlorine  are  the  constituents  of 
the  salt.  But  we  never  reverse  the  order  and  give  the  constituents 
first,  for,  if  we  did,  a  chemist  would  understand  us  to  mean  that 
metallic  sodium  and  yellow,  gaseous  chlorine  were  present  in  the 
tube,  which  would  not  be  the  case. 

When  a  body,  say  a  specimen  of  sulphur,  contains  a  little  of  some 
other  physical  component,  we  speak  of  it  as  impure  sulphur.  This 
does  not  mean  that  it  contains  dirt  in  the  ordinary  sense  of  the  term. 


INTRODUCTORY   II  23 

A  little  magnesium  chloride  is  a  common  impurity  in  table  salt 
(sodium  chloride),  and,  by  absorbing  moisture,  renders  it  more  moist 
in  damp  weather  than  it  would  otherwise  become.  Absolutely  pure 
bodies  are  unknown.  "  Chemically  pure  "  means  that  the  quantities 
of  the  impurities  which  the  material  is  most  apt  to  contain  have 
been  reduced  below  the  amount  which  would  interfere  with  the 
most  exact  chemical  work  for  which  the  substance  is  commonly 
employed. 

By  convention  we  continually  speak  of  "  pure  "  hydrochloric  acid, 
or  of  "  pure  "  sulphuric  acid,  although  there  may  be  60  per  cent  of 
water  present  in  the  former,  and  7  per  cent  in  the  latter.  By  this 
we  mean  to  distinguish  the  former  from  "  commercial  "  hydrochloric 
acid,  for  example,  which  contains,  in  addition  to  the  water,  impurities 
like  sulphuric  acid  and  a  coloring  matter.  The  water  is  in  fact 
disregarded,  since  it  is  assumed  to  be  present  in  all  cases. 

In  addition  to  the  foregoing,  there  are  three  other  sets  of  ideas, 
described  by  special  terms,  which  are  continually  used  in  chemistry. 

We  distinguish  one  body,  say  one  piece  of  sulphur,  from  another 
by  its  weight,  form,  or  volume.  Each  particular  specimen  differs 
from  every  other  in  these  attributes.  These  are  general  attributes 
possessed  by  matter  and  are  used  in  chemistry  for  measuring  quantity. 

When  all  the  bodies  to  which  we  should  apply  the  name  "  sul- 
phur "  are  compared,  we  find  that,  although  some  are  in  fine  powder, 
and  others  in  lumps  of  various  shapes  or  in  crystals,  and  thus  differ 
in  weight,  form,  and  volume,  they  nevertheless  have  many  qualities 
in  common.  These  qualities  we  call  specific  properties,  or  properties 
common  to  a  species.  The  material  composing  all  the  bodies  of  one 
species  we  call  a  substance.  Some  of  the  specific  properties  charac- 
terizing a  substance  and  common  to  all  specimens  of  one  species  are 
color,  odor,  crystalline  structure,  hardness,  melting-point  (tempera- 
ture of  fusion),  solubility  in  water  or  other  solvents,  boiling-point 
(temperature  above  which,  at  760  mm.  pressure,  the  substance  is 
gaseous),  specific  gravity,  specific  heat,  and  conductivity  for  elec- 
tricity. Thus,  sulphur  is  yellow,  has  little  odor,  crystallizes  in  the 
rhombic  system,  has  a  hardness  of  2.5  on  a  scale  of  ten,  has  the  m.-p. 
114.5°  C.,  is  not  perceptibly  soluble  in  water  but  dissolves  in  carbon 
disulphide  (41  parts  in  100  at  18°),  has  the  b.-p.  445°  C.,  the  sp.  gr.  2, 
the  sp.  ht.  0.18,  and  is  a  very  poor  conductor.  In  the  first  two 
chapters,  while  presenting  the  experimental  facts  required  for  our 


24  COLLEGE    CHEMISTRY 

discussion,  we  have  had  to  speak  of  substances  very  frequently  (e.g. 
pp.  2,  3,  and  5-8),*  and  have  done  so  always  in  the  above  sense. 

There  are  still  other  qualities  which  a  body  (or  specimen  of  mat- 
ter) may  possess.  It  has,  for  example,  a  certain  temperature,  press- 
ure, motion,  or  electric  charge.  These  are  impressed  upon  the  body 
and  do  not  inhere  in  it.  We  speak  of  them  as  conditions  of  the  body 
rather  than  properties.  In  the  use  of  the  body  they  may  be  altered, 
and  some  of  them  may  be  removed  or  added,  arbitrarily. 

Thus  there  are  three  kinds  of  qualities  to  be  considered.  The 
attributes,  like  weight,  form,  and  volume,  do  not  belong  to  substances 
but  do  belong  to  bodies.  The  specific  properties,  like  color,  solu- 
bility, and  odor,  belong  by  right  to  substances.  The  conditions,  like 
temperature  and  motion,  belong  to  neither,  for  they  can  be  altered 
without  changing  either  the  body  or  the  substance. 

Methods  of  Work  and  Observation  in  Chemistry. —  It  is  not 

the  end  of  chemical  work  to  make  generalizations  or  laws,  like  the 
characteristics  of  chemical  phenomena  (pp.  3,  11,  12, 18),  or  concep- 
tions, like  those  dealt  with  in  the  preceding  paragraph.  These  are 
simply  the  means  by  the  help  of  which  chemical  work,  whether  it 
be  investigation,  commercial  analysis,  or  manufacturing,  may  be 
carried  on  more  systematically.  Together  they  constitute  our 
system  for  classifying  the  facts  with  a  view  to  ready  reference.  The 
sample  experiments  (pp.  5-9),  if  reexamined,  will  show  that  we 
there  employed  most  of  the  categories  of  our  classification  which 
have  so  far  been  described. 

Thus,  in  the  experiment  with  iron  and  sulphur  (p.  6),  it  was  first 
our  object  to  find  out  whether  the  bodies  had  interacted  chemi- 
cally on  being  mixed.  To  do  this  we  noted  the  specific  properties 
(p.  23)*  of  the  substances  separately.  Using  these  properties  we 
were  able  to  identify  the  same  substances  in  the  mixture,  and  in 

*  References  to  previous  pages  are  used  in  order  to  save  needless  repetition 
in  writing.  But  the  beginner  requires  endless  repetition  in  his  reading  and 
must  form  the  habit  of  examining,  in  conjunction  with  the  current  text,  the 
parts  referred  to.  The  passages  cited  are,  by  the  reference,  made  part  of  the 
current  text,  which  will  usually  not  be  clear  without  them.  The  same  remark 
applies  to  topics  referred  to  by  name.  Such  topics,  if  treated  in  later  pages, 
are  distinguished  by  the  letters  q.v.,  and  must  be  sought  in  the  index. 

All  terms,  and  especially  those  borrowed  from  physics,  if  not  perfectly 
familiar,  must  be  looked  up  hi  a  work  on  physics  or  in  a  dictionary. 


INTRODUCTORY   II  25 

this  we  found  no  substance  with  new  properties.  Part  dissolved  in 
carbon  disulphide  and  reappeared  after  evaporation  of  the  solvent 
as  yellow,  rhombic  crystals,  and  the  rest  was  all  magnetic.  In  this 
connection  we  purposely  omitted  all  mention  of  the  quantity  and 
temperature,  because  attributes  (p.  23)  like  the  former  and  con- 
ditions (p.  24)  like  the  latter  do  not  characterize  substances  (since 
they  vary  with  each  specimen),  and  cannot  be  used  for  identification. 
While  all  the  specific  properties,  of  which  a  few  are  mentioned  on 
p.  23,  find  application  in  identification,  the  first  eight  on  the  list 
are  those  most  frequently  used.  Coincidence  in  two  or  three  specific 
properties  is  generally  sufficient  to  establish  identity. 

Most  of  the  properties  cannot  be  recognized  readily  in  mixtures, 
as  a  moment's  thought  will  show.  The  general  color  and  the  specific 
gravity  of  a  mixture,  containing  unknown  substances  in  unknown 
proportions,  for  example,  tell  us  little  about  the  corresponding 
properties  of  the  components.  Now,  there  are  few  pure  substances 
in  nature  or  in  the  products  of  experiment,  and  many  mixtures. 
Hence,  separation  of  the  components  of  a  mixture  usually  of  necessity 
precedes  the  process  of  identification  just  referred  to.  Thus,  we 
first  removed  the  sulphur  by  dissolving  it  and  then  recovered  it  by 
evaporation  of  the  solvent,  the  boiling-point  of  the  carbon  disulphide 
(46°)  being  much  lower  than  that  of  the  sulphur  (445°)  and  its 
vapor  pressure,  by  virtue  of  which  it  evaporated  at  the  temperature 
of  the  room,  being  therefore  relatively  high.  We  secured  the  iron 
because  of  its  insolubility.  Those  specific  properties  which  can  be 
used  for  separating  mixtures,  as  well  as  for  identification,  are  there- 
fore the  most  important  of  all.  They  are  the  melting-point  (see 
Sulphur  manufacture),  solubility,  boiling-point,  specific  gravity  (see 
Gold  extraction),  and  magnetic  qualities,  the  last  being  applicable 
almost  exclusively  to  iron,  however. 

In  connection  with  this  investigation  we  employed  several  of 
the  common  methods  of  manipulation  used  by  the  chemist.  These 
methods  are  derived  from  the  conceptions  described  in  the  last  para- 
graph. Thus  we  treated  the  mixture  with  a  solvent  (Fig.  2),  on  the 
assumption  that  if  it  was  heterogeneous  (p.  22)  the  components 
would  each  behave  as  if  alone  present.  We  then  filtered,  a  method 
invented  for  dealing  with  a  heterogeneous  mixture  consisting  of  a 
solid  and  a  liquid.  Decantation  is  often  used  in  such  cases  when  the 
solid  is  specifically  much  heavier  than  the  solvent  and  settles  readily. 


26 


COLLEGE    CHEMISTRY 


We  allowed  the  carbon  disulphide  to  evaporate  spontaneously,  and 
this  is  our  favorite  method  of  dealing  with  a  mixture  which  is  homo- 
geneous, and  therefore  would  run  through  a  filter  as  a  whole  without 
suffering  separation.  When  the  liquid  has  a  higher  boiling-point 
than  50-60°  C.,  as  water  has,  we  use  heat  from  a  steam-bath  or 
Bunsen  flame  to  promote  the  evaporation.  In  evaporation  we  allow 
the  vapor  of  the  liquid  to  escape,  because  it  is  the  less  volatile,  dis- 
solved body  that  we  wish  to  examine.  When  we  desire,  on  the  con- 
trary, to  examine  the  liquid,  the  vapor  must  be  condensed.  This 
method,  which  we  have  not  yet  had  occasion  formally  (see  Exercise 
2)  to  employ,  is  called  distillation  (Fig.  14).  The  jacket  round  the 


FIG.  14 

long  tube  is  filled  with  a  stream  of  cold  water,  which,  on  account  of 
its  high  specific  heat,  quickly  cools  and  condenses  the  vapor.  The 
resulting  liquid  is  caught  in  a  flask. 

These  methods  may  be  adapted  to  the  investigation  of  any  similar 
problem.  Thus,  gunpowder  is  made  by  the  intimate  mixing  of  sul- 
phur, charcoal,  and  saltpeter  (potassium  nitrate).  If  no  chemical 
interaction  whatever  has  occurred,  a  sample  will  be  wholly  separable 
into  these  components.  If  a  partial  change  has  taken  place,  a  certain 
amount  of  material  with  different  properties  will  be  discovered  in 
the  mixture.  If  the  change  has  been  complete,  no  portion  of  the 


INTRODUCTORY   II  27 

original  substances  will  be  found.  We  must  first  study  the  specific 
properties  of  each  of  the  ingredients  separately,  in  order  that  a  plan 
of  separation  may  be  devised,  and  that  we  may  have  a  basis  for 
comparison  with  the  products  of  the  separation. 

It  should  be  noted  that  our  methods  of  investigating  the  products 
of  nature  and  of  the  laboratory  are  purposely  limited  so  as  not  to 
separate  chemically  combined,  but  only  physically  mixed  forms  of 
matter.  After  a  physical  individual  has  been  isolated,  and  even  then 
only  if  it  has  new  properties,  and  is  not  recognized  as  a  known 
substance,  we  next  proceed  to  separate  it  into  its  chemical  constit- 
uents so  as  to  learn  which  constituents  it  contains  and  in  what 
relative  proportions  by  weight. 

The  experiments  with  mercuric  oxide  (p.  7)*  and  with  silver 
nitrate  (p.  8)  simply  rang  the  changes  on  the  same  conceptions, 
and  repeated  the  same  manipulations.  In  every  case  the  specific 
properties  —  color,  crystalline  form,  volatility,  solubility,  and  so 
forth  —  by  which  separation  and  recognition  were  effected  will  be 
found  to  have  been  mentioned. 

Summary.  —  In  this  chapter  we  have  added  considerably  to  our 
conception  of  the  scope  of  chemistry  (cf.  p.  11).  Although  our 
survey  is  by  no  means  yet  complete,  we  may  condense  our  results 
as  follows: 

Chemistry  deals  with  the  changes  in  composition  and  constitution 
which  substances  undergo  and  with  the  transformations  of  energy 
which  accompany  them.  To  convert  the  isolated  facts  into  a  science 
we  classify  related  parts  under  laws,  such  as  those  of  conservation 
of  mass  (p.  12)  and  of  energy  (p.  17),  and  under  conceptions,  such 
as  those  of  internal  energy  (p.  17),  element  (p.  20),  body  (p.  22), 
substance  (pp.  2-3,  21  and  23).  We  also  distinguish  between  attri- 
butes, specific  properties,  and  conditions  (pp.  23-24).  In  the  last 
paragraphs  we  have  indicated  briefly  the  use  to  which  these  concep- 
tions and  this  classification  are  put. 

The  system  of  classification  as  a  whole  is  part  of  the  everyday 
mode  of  thought  of  the  chemist,  for  thought  consists  largely  in  com- 
paring and  contrasting,  and  our  system  of  classification  furnishes 
the  plan  of  this  so  far  as  chemistry  is  concerned.  Learning  chemis- 
try consists,  therefore,  in  large  part,  in  learning  this  classification 
and  becoming  habituated  to  its  use. 

*  See  footnote  to  p.  24. 


28  COLLEGE   CHEMISTRY 

Exercises.*  —  1.  What  are  the  two  most  direct  ways  of  showing 
a  substance  to  be  a  compound?  Illustrate  each. 

2.  Discuss  in  detail  the  experiments  with  mercuric  oxide  (p.  7), 
and  with   silver  nitrate  (p.  8),  showing  what  specific  properties 
were  used  for  separating  and  identifying  the  products,  and  how  they 
answered  the  purpose.     Which  methods  of  manipulation  were  em- 
ployed in  the  second  experiment,  and  which   method  was  used, 
essentially,  in  the  first? 

3.  Define  the  following  terms,  and  find  illustrations  of  each,  other 
than  those  given  on  p.  22:  mixture,  physical  component,  chemical 
constituent. 

4.  Describe,  (a)  a  red-hot  rod  of  iron,  (b)  an  aqueous  solution  of 
sugar,  employing  all  the  terms  given  on  pp.  22-24,  so  far  as  they  are 
applicable. 

5.  What  (a)  elements  and  (b)  substances  are  contained  in  an 
aqueous  solution  of  sodium  nitrate?     Would  it  be  correct  to  say 
that  the  simple  substance  oxygen  is  contained  in  it?     What  then 
is  the  difference  in  meaning  between  the  terms  "  element  oxygen  " 
and  the  "  simple  substance  oxygen  "  ? 

*  The  exercises  should  in  all  cases  be  studied  with  minute  care.  They 
not  only  serve  as  tests  to  show  that  the  chapter  has  been  understood,  but 
very  frequently  (as  in  No.  1)  also  call  attention  to  ideas  which  might 
not  be  acquired  from  the  text  alone,  or  (as  in  No.  5)  assist  in  elucidating  ideas 
given  in  the  text  which,  without  the  exercises,  might  not  be  fully  grasped. 


CHAPTER  III 
INTRO  DUCTOBT   HI 

The  Law  of  Definite  Proportions,  Fifth  Characteristic  of 
Chemical  Phenomena.  —  The  ways  of  forming  or  decomposing  a 
compound,  or  of  carrying  out  a  more  complex  chemical  change,  may 
be  varied  indefinitely.  The  apparatus,  the  mode  of  experiment, 
and  the  proportions  of  the  materials,  may  be  altered  at  our  will. 
But,  in  spite  of  an  enormous  amount  of  careful  work,  no  case  of 
variation  in  the  proportion  of  the  constituents  actually  used  or  pro- 
duced in  a  given  chemical  action  has  come  to  light.  If  too  much  of 
one  constituent,  for  example,  is  taken,  a  part  simply  remains 
unchanged.  In  every  sample  of  each  compound  substance  formed  or 
decomposed,  the  proportion  by  weight  of  the  constituents  is  always  the 
same.  This  statement  of  fact  is  known  as  the  law  of  definite  propor- 
tions. Another  form  of  statement,  which  is  applicable  more  directly 
to  complex  chemical  actions,  is:  The  ratio  by  weight  of  any  one 
of  the  factors  or  products  of  a  chemical  change  to  any  other  is 
constant. 

The  Measurement  of  Combining  Proportions.  —  The  most 
exact  measurement  of  the  proportions  in  which  the  elements  combine 
to  form  compounds  involves  manipulations  too  elaborate  to  be  gone 
into  here.  One  or  two  brief  statements,  diagrammatic  rather  than 
accurate,  will  show  the  principles,  however. 

If  we  take  a  weighed  quantity  of  iron  in  a  test-tube  and  heat  it  with 
more  than  enough  sulphur  (an  excess  of  sulphur),  we  get  free  sulphur 
along  with  the  ferrous  sulphide  (pp.  6-7),  and  no  free  iron  survives. 
We  may  remove  the  free  sulphur  by  washing  the  solid  with  carbon 
disulphide.  The  difference  between  the  weights  of  ferrous  sulphide 
and  iron  gives  the  amount  of  sulphur  combined  with  the  known 
quantity  of  the  latter. 

As  an  example  of  the  study  of  rusting,  we  may  weigh  a  small 
amount  of  copper  in  the  form  of  powder  in  a  porcelain  boat  and  pass 


30 


COLLEGE   CHEMISTRY 


oxygen  over  the  heated  metal  (Fig.  15) .  If  we  limit  the  oxygen,  part 
of  the  copper  may  remain  unaltered;  if  we  use  it  freely,  the  excess 
will  pass  on  unchanged.  The  original  weight  of  the  copper,  and  the 


±11= 


Fia.  15 


increase  in  weight,  representing  oxygen,  give  us  the  data  for  deter- 
mining the  composition  of  cupric  oxide.  The  data  furnished  by 
one  rough  lecture-experiment,  for  example,  were  as  follows: 

Weight  of  boat  empty .  3.428  g. 

Weight  of  boat  +  copper 4.278  g. 

Difference  =  weight  of  copper 0.850 g. 

Weight  after  addition  of  oxygen 4.488g. 

Weight  without  oxygen 4.278  g. 

Difference  =  weight  of  oxygen 0.210  g. 

The  proportion  of  copper  to  oxygen,  so  far  as  this  one  measurement 
goes,  is  therefore  85  :  21. 

The  results  of  quantitative  experiments  are  usually  recorded  in  the 
form  of  parts  in  one  hundred.  To  find  the  percentage  of  each  con- 
stituent, we  observe  that  the  proportion  of  copper  is  85  :  85  +  21, 
or  -M-  of  the  whole.  That  of  the  oxygen  is  &fa  of  the  whole.  Thus 
the  percentages  are: 

Copper,         106  :  85  :  :  100  :  x        x   =  80.2 
Oxygen,        106  :  21  : :  100  :  x'      x*  =  19.8 


INTRODUCTORY   III  31 

Naturally,  the  mean  of  the  results  of  a  number  of  more  carefully 
managed  experiments  will  be  nearer  the  true  proportion.  The  per- 
centages at  present  accepted  as  most  accurate  are  79.9  and  20.1. 

In  the  case  of  mercuric  oxide,  we  may  decompose  a  known  weight 
of  the  oxide  (p.  7)  and  afterwards  weigh  the  mercury  and  ascer- 
tain the  oxygen  by  difference. 

The  following  figures  show  some  results  of  experiments  like  these, 
only  two  places  of  decimals  being  given.  The  numbers  in  paren- 
thesis will  be  explained  presently: 

(1)    Cupric  oxide  (2)  Mercuric  oxide 

Copper          79.9  F31.8 "]  Mercury        92.59  Tl 

Oxygen         20.1  L    8   J  Oxygen           7.41  L  8 

(3)   Water  (4)  Chlorine  monoxide 

Hydrogen     11.18  ri.008~|  Chlorine        81.59  r35.45"[ 

Oxygen         88.81  L    8    J  Oxygen         18.41  L    8    J 

Combining  and  Equivalent  Weights. —  The  percentages  in  the 
above  list  represent  the  true  proportions  by  weight  in  the  various 
compounds,  but  naturally  the  individual  numbers  constituting  those 
proportions  have  no  chemical  significance  whatever.  They  are  arbi- 
trary values  selected  so  that  the  constituents  of  the  proportion  may 
together  make  100.  Each  pair  represents  the  constant  ratio  which 
is  the  mean  result  of  numerous  experiments. 

We  begin  the  effort  to  reduce  these  numbers  to  order  by  selecting 
one  element  as  our  starting-point,  and  by  taking  some  convenient 
weight  of  it  as  the  basis.  As  we  shall  see,  it  makes  no  difference 
what  choice  we  make  in  either  respect.  To  avoid  waste  of  time,  we 
shall,  therefore,  ,use  oxygen,  as  it  is  the  element  generally  preferred 
by  chemists  for  the  purpose.  The  reason  for  this  preference  will 
be  apparent  later  (see  pp.  35  and  37). 

We  should  naturally  be  inclined  to  use  1  part  of  the  element  as  our 
basis.  But  our  later  steps  involve  finding  out  what  amounts  of  the 
other  elements  combine  with  this  quantity,  and  we  perceive  that  the 
amount  in  the  case  of  hydrogen  will  be  only  0.126  parts.  We  calcu- 
late this  from  (3):  88.81  :  11.18  :  1  :  x  (=  0.126).  If,  however,  we 
take  8  parts  of  oxygen,  this  amount  of  hydrogen  is  also  increased 
eight  times  and  becomes  1.008.  As  no  element  is  found  to  combine 
in  smaller  proportions  than  hydrogen,  we  are  satisfied  that  a  scale 


32  COLLEGE    CHEMISTRY 

for  our  numbers  based  on  8  parts  of  oxygen  will  not  involve  any 
values  less  than  1.  The  choice  of  scale  is  purely  one  of  con- 
venience. 

We  now  proceed  to  calculate  from  the  data  given  above  the  weight 
of  each  of  the  other  four  elements  which  combines  with  8  parts  of 
oxygen.  From  (1)  we  calculate  this  weight  of  copper,  thus,  20.1  : 
79.9  : :  8  :  x  (=31.8  parts  of  copper).  Similarly  we  find  that  8 
parts  of  oxygen  combine,  in  (2)  with  100  parts  of  mercury,  in  (3) 
with  1.008  parts  of  hydrogen,  and  in  (4)  with  35.45  parts  of  chlorine. 
Oxygen  unites  with  almost  all  the  known  elements,  and  these  four 
compounds  have  been  chosen  simply  as  a  sample  group. 

OXYGEN  COPPER  MERCURY  HYDROGEN  CHLORINE 

8  31.8  100  1.008  35.45 

Now  chlorine  combines,  not  only  with  oxygen,  but  also  with 
copper,  mercury,  and  hydrogen,  and  measurement  shows  that  the 
combining  proportions  are  represented,  exactly,  by  the  very  same 
numbers  as  before.  From  the  two  independent  facts  that  8  parts 
of  oxygen  combine  with  31.8  parts  of  copper  and  with  35.45  parts 
of  chlorine,  we  could  not  possibly  have  foretold  the  proportions  in 
which  copper  and  chlorine  would  combine  with  one  another.  Yet 
measurement  shows  it  to  be  31.8  :  35.45  precisely.  In  the  fol- 
lowing table  the  proportions  of  the  elements  are  placed  under  the 
corresponding  parts  of  the  names  of  the  compounds  with  chlorine: 

CUPRIC  CHLORIDE  MERCURIC  CHLORIDE  HYDROGEN  CHLORIDE 

31.8  :  35.45  100  :  35.45  1.008  :  35.45 

That  weight  of  each  element  which  combines  with  8  parts  of  oxygen 
(or  1.008  parts  of  hydrogen)  is  known  as  the  equivalent  weight  of  the 
element.  These  weights  are  equivalent  because  they  hold  identical 
amounts  of  oxygen  (or  hydrogen)  in  combination. 

If  additional  elements  had  been  included  in  the  group,  a  combining 
number  could  have  been  found  for  each,  and  seeming  coincidences  of 
the  same  nature  would  have  multiplied  rapidly.  Thus,  sulphur 
unites  with  hydrogen  to  give  hydrogen  sulphide.  If,  to  maintain  the 
same  scale,  we  use  1.008  parts  of  hydrogen  in  expressing  the  propor- 
tion, we  find  that  the  combining  ratio  is  1.008  :  16.03.  This  result 
could  not  enable  us  to  predict  the  proportion  of  copper  to  sulphur  in 


INTRODUCTORY   III  33 

cupric  sulphide,  but  measurement  shows  it  to  be  31.8  :  16.03.     And 
mercury  and  sulphur  unite  in  the  proportion  100  :  16.03. 

SULPHUR,  with  the  value  16.03,  may  therefore  be  added  to  the 
series  of  equivalent  weights. 

Law  of  Multiple  Proportions.  —  One  other  remarkable  fact 
remains  to  be  noted.  There  are  two  different  compounds  of  copper 
with  oxygen.  Cuprous  oxide  (q.v.),  the  one  not  mentioned  hitherto, 
is  found  in  nature  as  a  dark-red  mineral  which  is  entirely  different 
from  cupric  oxide  in  physical  properties.  It  can  also  be  prepared 
in  the  laboratory,  but  not  by  simply  passing  oxygen  over  heated 
copper.  Now,  analysis  shows  that  in  cuprous  oxide  the  proportion 
of  copper  combined  with  8  parts  of  oxygen  is  63.6.  This  new  number 
for  copper  is  not  unrelated  to  the  corresponding  value  in  cupric 
oxide  (viz.  31.8),  for  it  is  exactly  twice  as  great.  Again  mercury 
forms  mercurous  oxide  (see  p.  34)  as  well  as  mercuric  oxide  and  in 
the  former  the  proportion  of  oxygen  to  mercury  is  8  :  200.  The 
proportions  of  mercury  uniting  with  8  parts  of  oxygen  are  there- 
fore 100  and  200,  and  no  other  proportions  are  known.  Still  again, 
1.008  parts  of  hydrogen  unite  with  8  parts  of  oxygen  in  water  and 
with  2x8  parts  of  oxygen  in  hydrogen  peroxide.  The  fact  sug- 
gested by  these  three  examples  is  a  general  one.  It  was  discovered 
by  Dalton  (1804)  and  was  embodied  by  him  in  a  statement  known 
as  the  law  of  multiple  proportions,  which  ran  somewhat  as  follows: 
If  two  elements  unite  in  more  than  one  proportion  forming  two  or 
more  compounds,  the  quantities  of  one  of  the  elements,  which  in  the 
different  compounds  are  united  with  identical  amounts  of  the  other, 
stand  to  one  another  in  the  ratio  of  integral  numbers,  which  are  usually 
small. 


The  Law  of  Combining  Weights,  Sixth  Characteristic. — 

The  reader  should  now  examine  carefully  the  following  table  of 
combining  proportions.  It  includes  all  the  compounds  made  up  of 
pairs  of  the  six  sample  elements  under  consideration,  so  far  as  the 
existence  and  composition  of  such  compounds  have  been  deter- 
mined with  certainty.  The  substances  in  black  type  are  the 
ones  from  whose  composition  we  originally  derived  the  combining 
numbers,  the  others  illustrate  the  uniform  recurrence  of  the  same 
numbers: 


34 


COLLEGE    CHEMISTRY 


Gupric  oxide 
31.8  :  8.00 

Cuprous  oxide 

2  X  31.8  :  8.00 

Cupric  sulphide 

31.8  :  16.03 
Cuprous  sulphide 
2  X  31.8  :  16.03 

Cupric  chloride 

31.8  :  35.45 

Cuprous  chloride 
2  X  31.8  :  35.45 

Sulphur  dioxide 
16.03  :  2  X  8.00 

Sulphur  trioxide 
16.03  :  3  X  8.00 


Mercuric  oxide 
100.0  :  8.00 

Mercurous  oxide 

2  X  100.0  :  8.00 
Mercuric  sulphide 

100.0  :  16.03 
Mercurous  sulphide 

2  X  100.0  :  16.03 
Mercuric  chloride 

100.0  :  35.45 
Mercurous  chloride 

2  X  100.0  :  35.45 
Chlorine  monoxide 

35.45  :  8.00 
Chlorine  dioxide 

35.45  :  4  X  8.00 


Hydrogen  monoxide  (water) 

1.008  :  8.00 
Hydrogen  peroxide 

1.008  :  2  X  8.00 
Hydrogen  sulphide 

1.008  :  16.03 


Hydrogen  chloride 
1.008  :  35.45 


Sulphur  monochloride 
2  X  16.03  :  35.45 

Sulphur  tetrachloride 
16.03  :  2  X  35.45 


It  will  be  observed  that  the  same  number  serves  always  to  express 
the  combining  proportions  of  a  given  element,  provided  that  multi- 
plication by  an  integer  is  permitted  when  necessary.  There  are  also 
some  combinations  of  three  of  the  same  elements  in  one  compound. 
But  even  in  such  cases  the  fundamental  numbers  still  reappear. 
Thus  oxygen,  chlorine,  and  hydrogen  combine  in  the  proportions: 
2X8:  35.45  :  1.008,  6x8:  35.45  : 1.008,  and  8  X  8  :  35.45  :  1.008. 
Nor  is  the  recurrence  of  a  fundamental  number  an  exclusive  property 
of  the  six  elements  we  have  chosen  for  illustration.  A  vast  table 
containing  every  known  element,  and  every  compound  of  every  ele- 
ment, if  prepared  in  the  same  way  as  that  given  above  (by  starting 
with  a  fixed  number  for  one  element  and  calculating  the  combining 
proportions  in  all  the  compounds  with  this  number  as  a  basis),  would 
show  that  every  element  used  an  individual,  fundamental  number, 
or  integral  multiples  of  this  number,  in  every  one  of  its  combinations. 
No  exceptions  to  this  rule  would  be  found  in  the  whole  mass  of  data. 

The  law  of  combining  weights  may  be  put  briefly  thus:  The  propor- 
tions by  weight  in  which  all  chemical  combinations  take  place  can  be 
expressed  in  terms  of  small  integral  multiples  of  fixed  numbers,  called 
combining  weights,  one  for  each  element.  It  describes  what  is  perhaps 
the  most  striking  of  all  the  characteristics  of  chemical  action. 

If  the  reader  will  now  reexamine  carefully  the  way  in  which  the 


INTRODUCTORY    III  35 

data  were  handled,  the  following  significant  facts  will  be  noted: 
(1)  Oxygen  was  made  the  starting-point  of  the  system  and  the  value 
8  was  assigned  to  its  combining  weight.*  Had  a  different  value 
been  used,  all  the  numbers  would  have  been  different.  But  the 
change  would  have  affected  all  the  numbers  in  the  same  proportion 
and  so  only  the  scale  of  the  numbers  would  have  been  affected.  An 
individual  combining  number,  one  for  each  element,  would  have 
still  recurred  wherever  the  element  itself  appeared.  (2)  Even  use 
of  another  element  as  the  initial  one  would  not  have  prevented  the 
discovery  of  the  law.  Thus  with  hydrogen  as  the  initial  element, 
and  the  value  1.008,  no  change  whatever  would  have  been  noted. 
With  a  different  value  for  hydrogen,  a  change  of  scale  in  all  the  num- 
bers would  have  followed,  but  individual  combining  weights  would 
have  appeared  as  before.  (3)  Finally,  if  cuprous  oxide  had  been  used 
instead  of  cupric  oxide  for  the  first  proportion,  the  value  found  for 
copper  would  have  changed  to  63.6,  while  the  other  numbers  would 
have  remained  unaffected.  But  this  number  would  serve  as  the 
combining  weight  of  copper  just  as  well  as  31.8,  for  the  composition 
of  cupric  oxide  can  be  expressed  by  the  proportion  63.6  :  2  X  8 
as  well  as  by  31.8  :  8,  and  that  of  cupric  sulphide  by  63.6  :  2  X  16.03 
as  well  as  by  31.8  :  16.03.  An  important  conclusion  therefore  fol- 
lows from  these  considerations,  and  we  shall  have  occasion  to  use  it 
presently,  namely,  that  any  one  or  more  of  the  equivalent  weights 
(with  scale  oxygen  =  8)  may  be  separately  multiplied  by  an  integer 
without  its  usefulness  as  a  member  of  the  series  being  at  all  impaired. 
The  importance  of  the  fact  described  in  the  law  of  combining 
weights  cannot  be  emphasized  too  strongly.  Without  this  fact,  the 
remembering  of  the  compositions  of  chemical  substances,  necessary 
as  it  is  to  the  chemist,  would  have  been  completely  beyond  the  power 
of  any  ordinary  memory.  With  it,  the  task  becomes  comparatively 

*  Oxygen  was  chosen  as  the  basis  of  the  system  because  the  exact  deter- 
minations of  the  combining  weights  of  most  of  the  elements  have  actually 
been  made  by  direct  union  with  oxygen  or  with  the  help  of  but  one  inter- 
mediate step.  If  the  question  had  been  one  of  mathematics,  hydrogen,  the 
element  with  the  lowest  combining  proportions,  would  have  furnished  the 
basis  and  unit  of  the  scale.  But  the  question  was  the  practical  one  of  getting 
the  most  accurate  measurements  for  the  relative  magnitudes  of  the  numbers, 
so  oxygen  was  chosen  instead.  Nevertheless,  the  value  8  was  selected  in 
order  that  the  advantage  of  having  a  mathematical  unit,  or  something  close 
to  it,  in  the  combining  weight  of  hydrogen,  might  be  retained  also. 


36 


COLLEGE    CHEMISTRY 


simple.  It  is  only  necessary  to  decide  on  the  best  system  of  values 
for  the  combining  weights,  and  then,  regarding  the  value  of  this  for 
each  element  as  the  unit  of  weight  for  that  element,  to  express  the  pro- 
portions of  the  element  in  every  compound  by  the  proper  multiples. 
Thus,  given  a  list  of  the  combining  weights,  one  for  each  element,  only 
the  small  integral  multiples  have  to  be  kept  in  mind  in  connection 
with  each  compound. 

The  reader  will  require  a  little  time,  however,  before  he  becomes 
accustomed  to  the  use,  not  of  a  single  unit  of  weight,  but  of  a  different 
one  for  each  element.  Chemistry  is  the  only  science  in  which  the 
physical  unit  of  weight,  which  is  the  same  for  all  materials,  is  not 
employed  for  every  purpose.  The  physical  manipulations  of  the 
chemist  are  carried  out  with  the  use  of  physical  units,  but  the  chemi- 
cal results  are  expressed  in  terms  of  individual  unit  quantities  of  the 
several  elements,  the  combining  weights. 

The  individual  units  actually  used  for  each  element  are  not  in  all 
cases  identical  with  those  we  have  given.  The  final  values  will  be 
discussed  in  the  next  section. 

Atomic  Weights.  —  The  chemist  frequently  uses  the  idea  of 
equivalents  and  the  values  we  have  given  them.  But  far  more  often 
he  employs  a  slightly  differing  set  of  numbers,  which,  for  reasons  that 
will  appear  in  a  subsequent  chapter,  he  calls  atomic  weights.  The 
following  list  shows  the  elements  whose  equivalents  we  have  been 
discussing,  along  with  one  or  two  others,  added  by  way  of  furnish- 
ing a  fair  sample,  and  gives  both  sets  of  weights  for  the  purpose  of 
comparison : 


Equiva- 

Equiva- 

Element. 

lents  (Com- 
bining 

Atomic 
Weights. 

Element. 

lents  (Com- 
bining 

Atomic 
Weights. 

Weights). 

Weights). 

Oxvffen  . 

8 

16 

Iron     

27.95 

55.9 

Copper   .... 
Sulphur     .   .    . 

31.8 
16.03 

63.6 
32.06 

Magnesium    . 
Carbon   .    .    . 

12.18 
3.00 

24.36 
12.00 

Mercury.    .    .    . 

100.0 

200.0 

Aluminium    . 

9.03 

27.1 

Chlorine     .    .    . 

35.45 

35.45 

Sodium  .    .    . 

23.05 

23.05 

Hydrogen  .    .    . 

1.008 

1.008 

Bromine     .    . 

79.96 

79.96 

It  will  be  seen  that  some  equivalents  have  been  multiplied  by  two, 
the  first  four  and  those  of  iron  and  magnesium,  for  example;  some 


INTRODUCTORY   III  37 

have  been  multiplied  by  three,  like  that  of  aluminium;  some  by  four, 
like  that  of  carbon;  and  some  remain  unchanged,  like  those  of 
chlorine,  hydrogen,  sodium,  and  bromine. 

The  reasons  for  this  manipulation  of  the  simple  equivalents  found 
by  experiment  are  based  upon  theoretical  considerations.  As  it  is 
impossible,  until  we  reach  certain  important  facts  which  cannot  be 
introduced  here,  to  explain  these  considerations,  the  discussion  of 
the  reasons  for  the  changing  of  the  numbers  will  be  postponed  until 
after  these  facts  are  before  us.  Suffice  it  to  say  that  great  advan- 
tages are  found  to  attach  to  these  modifications  in  the  values.  The 
step  from  equivalents  to  atomic  weights  (q.v.)  is  taken  before  the 
justification  of  it  can  be  given,  because  otherwise  formulae  (see  next 
chapter)  could  not  be  used  in  the  earlier  chapters,  and  so  the  advan- 
tages their  employment  offers  would  be  sacrificed. 

The  atomic  weight  is  the  unit  of  weight  (p.  36)  actually  used  in 
expressing  the  proportions  of  each  element  in  all  its  compounds. 
The  integral  factors  are,  of  course,  different  from  those  which  would 
be  employed  in  expressing  the  composition  of  the  same  substance  in 
terms  of  equivalents,  because  many  of  the  latter  have  been  multi- 
plied by  small  integers  already  in  course  of  being  made  into  atomic 
weights.  But  the  multiplication  has  in  every  case  been  by  an  integer, 
so  that  the  new  numbers  are  just  as  serviceable  as  the  old  (p.  35). 

To  the  reasons  given  above  for  the  choice  of  oxygen  as  the  funda- 
mental element,  and  the  value  8  for  its  combining  or  equivalent 
weight,  one  other  may  now  be  added.  The  majority  of  the  atomic 
weights,  calculated  on  this  basis  from  the  experimental  results,  fall 
so  close  to  being  integers  that  the  nearest  round  numbers  are  exact 
enough  for  ordinary  use.  Thus  in  the  above  list  nine  of  the  twelve 
atomic  weights  are  within  0.1  of  the  nearest  whole  number.  This 
convenience  disappears  when,  for  example,  hydrogen  with  the  value 
1  is  made  the  basis. 

The  reader  will  inevitably  find  difficulty  at  first  in  thoroughly 
grasping  the  significance  of  these  numbers.  It  may,  therefore,  be  of 
some  assistance  if  a  hint  is  thrown  out  which  will  suggest  a  concrete 
basis  for  this  curious  property.  These  numbers  appear  to  mean 
that,  when  we  wish  to  make  a  chemical  compound,  we  may  choose 
any  two  elements  from  the  list,  and,  if  it  is  found  that  they  can  com- 
bine at  all,  we  have  only  to  take  the  atomic  weights,  worked  out 
from  other  combinations  of  each  element,  and  we  shall  find  that  they 


38  .  COLLEGE   CHEMISTRY 

will  exactly  suffice  for  this  case  of  chemical  union.  If  complete  com- 
bination of  both  materials  does  not  take  place,  then  trial  will  quickly 
show  what  multiples  of  the  atomic  weights  will  result  in  this.  The 
situation  seems  to  suggest  that  the  constructing  of  chemical  com- 
pounds depends  upon  the  putting  together  of  ready-made  "  parts," 
like  those  of  a  watch  or  a  bicycle.  The  parts  seem  to  be  "  inter- 
changeable," and  each  element  seems  to  be  furnished  to  us  by  nature 
in  ready-made  packets  suitable  for  application  in  building  up  any 
chemical  structure. 

A  complete  list  of  atomic  weights  is  printed  ori  the  inside  of  the 
cover  at  the  back  of  this  book. 

Summary.  —  This  chapter  adds  an  important  item  to  our  state- 
ment of  the  scope  of  the  science  (cf.  p.  27),  which,  therefore,  now 
reads  as  follows:  Chemistry  deals  with  the  quantitative  study  of  the 
changes  in  composition  and  constitution  which  substances  undergo 
and  with  the  transformations  of  energy  which  accompany  them.  To 
express  the  quantitative  relations  which  are  observed,  a  different  unit 
of  weight  is  employed  for  each  element,  and  is  known  as  the  atomic 
weight  of  the  element. 

Exercises. —  1.  To  test  the  correctness  of  the  statements  on 
p.  35,  take  mercury  as  the  basal  element  and  250.  as  its  combining 
weight,  and  work  out  from  the  data  on  pp.  31  and  32  (second  line 
from  the  foot)  the  corresponding  combining  weights  of  the  other  five 
elements.  Then  show  that  the  values  obtained  have  the  same 
property  as  have  the  equivalents. 

2.  Express  in  terms  of  atomic  weights,  or  their  integral  multiples, 
the  composition  of  cupric  oxide,   cupric  chloride,  sulphur  mono- 
chloride. 

3.  Show  that  doubling  the  atomic  weight  of  chlorine  would  give 
an  available  combining  number. 


CHAPTER  IV 
INTRODUCTORY  IV 

A  CONSIDERATION  of  the  contents  of  the  foregoing  chapter  will 
show  that  the  complete  description  of  a  chemical  change  must  be 
exceedingly  involved.  In  a  moderately  complex  action,  such  as 
that  of  sodium  chloride  upon  silver  nitrate,  we  should  say  that  sodium 
chloride,  composed  of  one  atomic  weight  each  of  sodium  and  chlorine, 
when  brought  in  contact  with  silver  nitrate,  composed  of  one  atomic 
weight  each  of  silver  and  nitrogen  and  three  atomic  weights  of  oxy- 
gen, gave  silver  chloride,  composed  of  one  atomic  weight  each  of  silver 
and  chlorine,  and  sodium  nitrate,  composed  of  one  atomic  weight 
each  of  sodium  and  nitrogen  and  three  atomic  weights  of  oxygen. 
Such  a  statement,  while  it  would  give  all  the  facts  in  the  quantitative 
point  of  view,  would  be  difficult  to  grasp  and  lacking  in  perspicuity. 

Symbols,  Formulae,  and  Equations.  —  In  order  to  represent 
the  nature  of  a  chemical  change  in  a  form  which  may  be  taken  in  at 
a  glance,  the  chemist  is  in  the  habit  of  using  certain  symbols,  first 
introduced  by  Berzelius.  Thus,  the  letters  Ag  represent  one  atomic 
weight  (i.e.,  107.93  parts)  of  silver  (argentum) ,  and  O  represents  one 
atomic  weight  (i.e.,  16  parts)  of  oxygen.  In  other  words,  the  symbol 
of  an  element  means  one  chemical  unit  weight  of  the  element.  Since 
many  elements  begin  with  the  same  initial,  two  letters  have  fre- 
quently to  be  used  to  distinguish  them.  When  the  names  of  the 
elements  are  not  the  same  in  all  languages,  resort  is  frequently  had 
to  Latin.  Thus,  Cu  stands  for  one  combining  weight  of  copper 
(cuprum),  Fe  is  used  for  iron  (ferrum),  Hg  for  mercury  (hydrargyrum). 
From  German  we  have  Na  for  sodium  (natrium)  and  K  for  potassium 
(kalium).  To  represent  a  compound,  the  symbols  of  the  elements 
which  it  contains  are  placed  side  by  side,  small  numbers  indicating 
multiples  of  the  atomic  weights  where  they  occur.  Thus,  sodium 
chloride  is  represented  by  the  expression  NaCl,  silver  nitrate  by 
the  symbols  AgNO3.  A  combination  of  symbols  is  called  a  formula. 


40  COLLEGE    CHEMISTRY 

The  symbols  composing  a  formula,  taken  by  themselves,  do  not 
stand  for  any  definite  quantity;  each  is  one  factor  of  a  proportion. 
Ag  means  the  proportion  of  107.93  parts  of  silver  to  the  proportions 
of  the  other  elements  represented  by  the  other  symbols  which  may 
be  connected  with  it. 

The  value  of  these  symbols  lies,  not  only  in  the  compact  way  in 
which  the  resulting  formulae  present  the  composition  of  compounds, 
but  also  in  the  use  which  may  be  made  of  them  in  showing  at  a  glance 
the  details  of  a  chemical  change.  The  chemical  action  just  mentioned 
appears  as  follows: 

NaCl  +  AgNO3  -» AgCl  +  NaN03. 

This  expression  contains  all  that  was  conveyed  by  the  words  which 
were  written  out  in  full  above.  The  arrow  indicates  that  the  mate- 
rials on  the  left-hand  side  pass,  in  the  chemical  transformation,  into 
those  on  the  right-hand  side.  Such  symbolic  expressions  are  called 
equations. 

One  other  variation  is  in  frequent  use.  The  equation  for  the 
formation  of  rust,  if  we  left  the  water  which  enters  into  its  com- 
position out  of  the  question,  would  run  thus: 

2Fe  +  3O  ->  FeA. 

It  will  be  observed  that  we  employ  the  form  2Fe  before  combination 
and  Fe2  (in  Fe2O3)  after  it.  The  reasons  for  this  usage  will  become 
clear  as  we  proceed.  ^Ve  note  simply  that  2Fe  means  2  separate 
atomic  weights  of  iron,  as  3Fe2O3  would  mean  three  separate  formula- 
weights  of  oxide  of  iron.  The  same  substance  (cf.  p.  21),  iron,  might 
appear  as  5Fe  or  8Fe  in  other  equations,  according  to  the  proportion 
needed.  But  Fe203  is  a  group  of  five  atomic  weights  united  chemi- 
cally. The  substance  ferric  oxide  never  contains  any  more  than 
two  aj-omic  weights  of  the  element  iron,  and  its  formula  is  invariable. 
Thus  the  regular  integers  multiplying  the  atomic  weights  in  the  com- 
position of  a  particular  compound  are  written  after  the  symbols  of 
the  elements,  while  more  arbitrary  coefficients  which  change  from 
one  use  of  the  substance  to  another  are  written  in  front.  When  no 
coefficient  appears  in  front  of  a  symbol  or  formula,  1  is  to  be  under- 
stood. 

\l    Making  Formulce.  —  To  make  the  formula  of  a  compound  sub- 
stance, assuming  the  formula  to  be  unknown,  two  kinds  of  informa- 


INTRODUCTORY   IV  41 

tion  are  required.  We  ascertain  (1)  by  measurement  the  proportion 
by  weight  of  the  constituents  in  the  compound.  We  require  also 
(2)  to  know  the  chemical  unit  weights  —  the  atomic  weights  — 
which  have  been  accepted  by  chemists  for  each  constituent  element. 
By  factoring  the  terms  of  the  first  proportion  so  that  one  factor  in 
each  case  is  the  atomic  weight,  we  discover  whether  multiples  of 
the  atomic  weights  will  be  required  to  represent  the  composition  of 
the  substance,  and  if  so  what  these  must  be.  An  illustration  will 
make  the  process  clear. 

Suppose  the  problem  is  to  make  the  formula  of  dried  rust.  By 
weighing  before  and  after  the  change,  we  get  the  weight  of  the  iron 
and  of  the  corresponding  amount  of  oxygen  in  the  rust  it  produces. 
If  we  took  2  g.  of  iron  we  should  get  about  2.86  g.  of  rust.  So  that 

2 

the  proportion  of  iron  to  oxygen  is .     Now,  in   the   formula, 

0.86 

the  same  ratio  must  be  represented  by  means  of  multiples  of  the 
atomic  weights  (p.  37).  We  therefore  divide  the  quantity  of  each 
element  by  the  corresponding  atomic  weight.  This  gives  us  the 
factors  by  which  the  atomic  weights  are  to  be  multiplied.  The 
atomic  weights  are  55.9  and  16  respectively:  2  -5-  55.9  =  0.0358, 

2 
and    0.86  -s-  16  =  0.0537.      The    proportion   --^   then    becomes 

U.oO 

55.9  X  0.0358  ^  NQW  ^^  prOpOrtiOn  must  be  capable  of  expres- 
16.0  X  0.0537 

sion  in  terms  of  integral  multiples  of  the  atomic  weights.  We 
find  that  the  greatest  common  measure  of  the  two  factors  is 
0.0179.  Dividing  above  and  below  by  this,  we  obtain  the  ratio 

X      .      Substituting  the  symbols  for  the  atomic  weights,  the 
lo.O  X  o 

proportion  appears  as    **  * — ,    and  the  formula  is  therefore  Fe2O3. 

O  X  o 

It  is  obvious  that  setting  the  symbols  down  side  by  side  is  not 
sufficient.  We  must  determine  by  measurement  the  factors  by 
which  they  are  to  be  multiplied. 

Making  Equations.  —  To  make  the  equation  representing  a 
chemical  change  we  note,  (1)  what  substances  were  used,  and  ascer- 
tain, by  study  of  their  properties,  what  substances  were  formed. 
Then  we  learn  (2)  the  formulae  of  the  substances  used  and  pro- 


42  COLLEGE   CHEMISTRY 

duced.  This  we  do  either  by  measurement  and  calculation,  as 
shown  above,  or  we  find  in  the  text  the  formulae  as  they  have  been 
determined  by  the  experimental  work  of  chemists.  From  these  we 
prepare  (3)  a  skeleton  equation  which,  in  the  instance  discusseci, 
would  appear  thus: 

Fe  -f  0 


We  are  careful  to  place  the  initial  substances  on  the  left,  and  to 
point  the  arrow  towards  those  which  are  produced.  Finally  (4) 
we  "  balance  the  equation  "  by  placing  the  proper  coefficients  before 
the  formulae.  This  last  operation  requires  experience  for  its  rapid 
performance.  A  good  rule  is  to  begin  by  picking  out  that  one  of  the 
formulae  which  contains  the  largest  number  of  atomic  weights,  no 
matter  upon  which  side  it  appears.  Here,  this  formula  is  Fe2O3. 
We  then  reason  that,  to  obtain  F^  we  require  2Fe,  and  to  obtain 
O3  we  require  3O,  and  accordingly  we  place  these  coefficients  before 
the  appropriate  symbols,  thus: 

2Fe  +  30  ->  FeA- 

It  is  hardly  necessary  to  add  that  a  chemical  equation  gives  the 
proportions  of  the  materials  and  nothing  more.  The  physical  condi- 
tions, for  example,  whether  the  substances  are  dissolved  in  a  liquid, 
or  are  in  the  state  of  gas,  or  are  at  a  high  temperature,  have  no  place 
in  it.  The  physical  properties  of  the  substances  concerned,  and  also 
the  energy  in  the  form  of  heat  or  electricity  which  may  appear  or 
disappear  in  the  process,  are  likewise  left  entirely  out.  A  question 
in  regard  to  the  nature  of  a  particular  chemical  change  demands  in 
answer  a  full  statement  of  all  these  things.  The  equation  is  therefore 
an  essential  part,  but  only  a  part,  of  such  a  statement. 

Units  of  Measurement  in  Chemical  Work.  —  In  chemical 
work  temperatures  are  invariably  measured  on  the  Centigrade  scale. 
The  temperature  of  a  mixture  of  ice  and  water  is  the  zero  point. 
The  temperature  of  the  steam  which  rises  from  water  boiling  under 
a  pressure  of  one  atmosphere  is  represented  by  100°.  The  interval 
between  those  two  points  is  divided  into  one  hundred  equal  parts. 

For  the  expression  of  length,  weight,  and  volume,  the  metric  sys- 
tem is  employed.  The  unit  of  this  system  is  the  meter,  which  is  sub- 
divided into  decimeters,  centimeters  (cm.),  and  millimeters  (mm.). 


INTRODUCTORY   IV  43 

For  small  measurements  the  last  subdivision  is  taken  as  the  unit. 
A  cubic  centimeter  (c.c.)  is  the  unit  of  volume  for  small  measure- 
ments. For  larger  ones  the  liter,  which  contains  1000  cubic  centi- 
meters, is  used.  The  unit  of  weight  is  that  of  one  cubic  centimeter 
of  water  at  4°,  the  temperature  of  maximum  density.  This  is  called 
the  gram.*  For  larger  amounts  of  material  the  kilogram,  which 
contains  1000  grams  (1000  g.),  is  frequently  employed.  The  meter 
is  equal  to  about  39J  inches  in  ordinary  measures,  and  the  centi- 
meter is  very  nearly  £  of  an  inch.  One  liter  is  about  ^  of  a  cubic 
foot  and  contains  61  cubic  inches  or  35  fluid  ounces.  One  hundred 
grams  is  about  3J  ounces  avoirdupois,  and  one  ounce  equals  28.35 
grams. 

Calculations  in  Chemistry.  —  In  the  laboratory  it  is  frequently 
desirable  that  we  should  know  what  amount  of  some  substance  may 
be  obtained  by  a  given  chemical  action  from  another,  or  what 
amount  of  material  must  be  used  to  obtain  the  desired  amount  of 
some  product.  This  information  is  readily  accessible,  since  measure- 
ments of  quantity  in  connection  with  most  chemical  changes  are  on 
record.  The  simplest  and  most  easily  handled  form  of  this  record 
is  found  in  the  formulae  of  compounds,  and  in  the  equations  repre- 
senting the  changes  which  they  undergo.  It  is  most  convenient, 
therefore,  when  a  question  of  this  kind  occurs,  to  ascertain  and  write 
down,  first,  the  equation.  Having  then  before  us  the  information  in 
regard  to  the  quantities  in  the  most  condensed  form,  we  may  use 
such  parts  of  this  information  as  are  required  for  the  problem  in  hand. 

Suppose,  for  example,  that  the  question  is  in  regard  to  the  weight 
of  oxygen  which  may  be  obtained  from  120  g.  of  mercuric  oxide.  We 
write  down  the  equation,  and,  if  the  numbers  are  not  familiar  to  us, 
we  ascertain  the  atomic  weights,  a  table  of  which  is  printed  on  the 
inside  of  the  rear  cover  of  this  book.  These  we  place  then  below  the 
symbols  by  which  they  are  represented,  thus: 

HgO  -»  Hg  +  O 

2004-16  200         16 

Reading  this  equation,  it  appears  that  one  formula- weight,  or  216 
parts  by  weight,  of  mercuric  oxide  give  one  atomic  weight,  16  parts, 

*  In  point  of  fact,  the  gram  is  the  one-thousandth  part  of  the  weight  of  the 
standard  kilogram  kept  in  Paris.  This  differs  from  the  weight  of  1  c.c.  of  water 
at  4°  by  less  than  0,01  per  cent. 


44  COLLEGE   CHEMISTRY 

of  oxygen,  and  the  question  is,  What  weight  of  oxygen  will  bi 
obtained  from  120  grams  of  the  oxide?  The  answer  may  be  stated 

by  simple  proportion:  216  :  16  :  :  120  :  x,  or  16  X  —  =  Answer. 

The  reader  must  conquer  a  tendency  to  speak  of  the  symbol  0  as  repre- 
senting "  1  part  "  of  oxygen:  it  stands  for  16  parts.  The  word  "  part  " 
refers  to  physical  units  exclusively. 

It  will  be  noticed  that  not  all  the  data  which  we  have  written 
down  are  necessarily  used.     In  general,  only  two  of  the  three  or 
more  weights  which  the  equation  represents  will  be  required. 
2L(  £,  "X£ro  .'    (W  •  7^ 

Exercises. —  1.  What  weight  of  mercury  is  obtained  from  120  g. 
of  mercuric  oxide? 

2.  What  weight  of  mercuric  oxide  will  furnish  20  g.  of  oxygen?  „ 

3.  What  weight  of  rust  may  be  obtained  from  10  g.  of  oxygen? 

4.  What  weight  of  silver  chloride  is  obtained  from  50  g.  of  silver 
nitrate  (p.  40)? 

5.  How  much  silver  is  contained  in  100  g.  of  an  impure  specimen 
of  silver  chloride  which  is  33  per  cent  sand? 

6.  If  26  g.  of  mercurous  oxide  are  required  to  give,  by  heating, 
1  g.  of  oxygen,  what  is  the  formula  of  the  substance? 

7.  What  are  the  formulae  of  the  substances  possessing  the  follow- 
|    ing  percentage  compositions?    The  percentages  are  to  be  divided  by 

^  the  atomic  weights  just  as  were  the  actual  weights  in  the  illustration 
p.  41. 

I  X    II  III 

Magnesium,      25.57     Sddium,     32.43     Potassium,    26.585 
Chlorine,  74.43     Sulphur,    22.55     Chromium,  35.390 

Oxygen,    45.02     Oxygen,        38.025 

"  8.  What  are  the  percentage  compositions  of  substances  possessing 
the  following  formulas :  Mn304,  KBr,  Fe§04? 

9.  Compare  the  formula  of  mercurous  oxide,  found  in  6,  with  that 
of  mercuric  oxide,  and  show  how  the  compounds  illustrate  the  law 
of  multiple  proportions  (p.  33). 

10.  Using  the  proportions  given  on  p.  34,  calculate  the  formulae 
of  all  the  compounds  in  the  table.     Compare  in  each  case  the  factors 
which  multiply  the  symbols  (and  therefore  the  atomic  weights) 
with  those  used  in  the  table  to  multiply  the  equivalent  weights. 
Explain  the  differences  where  they  occur. 

-rvu-  .    r.i'v^-M 


j 


• 


CHAPTER 


OXYGEN 

Historical  and  Introductory.  —  Almost  onie-quarter  of  the 
atmosphere,  by  weight,  is  free  oxygen.  Water  contains  nearly  89 
per  cent  of  oxygen  in  combination,  and  this  element  constitutes 
about  50  per  cent  of  common  materials  like  sandstone,  limestone, 
brick,  and  mortar.  On  account  of  its  predominance  over  other 
elements  in  quantity  (p.  21),  and  the  exceptional  capacity  which 
it  exhibits  for  forming  compounds  with  a  great  variety  of  other 
elements,  the  systematic  study  of  chemistry  may  conveniently  be 
begun  with  oxygen. 

While  many  elements  which  are  less  easily  obtainable  than  oxygen 
have  been  recognized  as  distinct  substances  for  many  centuries, 
oxygen  did  not  attain  this  position  until  it  was  prepared,  first  by 
Scheele  in  Sweden  in  1771-3  and,  independently,  by  Priestley  in 
1774.  The  latter  was  particularly  interested  in  examining  the 
nature  of  the  gases  which  were  evolved  by  some  materials  when 
heated.  He  found  that  mercuric  oxide  gave  off  an  unusual  amount 
of  a  gas,  or  "  air  "  as  he  called  it.  Priestley  discovered  that  this 
gas  supported  combustion  extremely  well  and,  later,  that  it  was 
respirable  and  favorable  to  the  life  of  small  animals,  such  as  mice. 
He  did  not,  however,  recognize  that  atmospheric  air  was  a  mixture, 
and  that  the  substance  he  had  obtained  was  in  reality  identical  with 
that  component  of  the  air  which  has  the  same  properties. 

Meanwhile  Lavoisier,  in  Paris,  who  had  been  studying  the  rusting 
of  metals  in  the  air,  heard  of  Priestley's  experiments,  and  demon- 
strated that  the  latter's  "  good  air  "  was  really  a  component  of 
common  air,  and  combined  with  metals  when  they  formed  rusts,  or 
"  calces,"  as  they  were  then  called.  He  proved  this  conclusively 
by  heating  mercury  in  an  inclosed  volume  of  air.  The  red  mercuric 
oxide  accumulated  on  the  surface  of  the  mercury  and  simultaneously 
the  air  suffered  a  shrinkage  of  a^out  one-fifth  of  its  volume.  The 
residual  gas,  nitrogen,  no  longef 'supported  combustion  or  life.  The 

45 


'&> 


46  COLLEGE    CHEMISTRY 

oxide  on  being  heated  more  strongly  by  itself  gave  off  a  gas  whose 
volume  exactly  corresponded  with  the  shrinkage  undergone  by  the 
inclosed  air,  and  this  gas  possessed  in  an  exaggerated  degree  the 
properties  which  the  air  had  lost.  The  proof  that  oxygen  was  a  con- 
stituent of  the  atmosphere  was  therefore  complete.  Lavoisier,  in 
the  mistaken  belief  that  the  new  element  was  an  essential  con- 
stituent of  all  sour  substances,  named  it  oxygen,  or  acid-producer 
(Gk.  o£vs  an  acid,  yewav  to  produce). 


J 


Preparation  of  Oxygen.  —  1.  Oxygen  may  be  separated  from 
the  other  substances  mixed  with  it  in  the  atmosphere  by  liquefying 
the  air  (see  Liquid  air),  allowing  the  nitrogen,  which  is  more  volatile, 
to  escape,  and  finally  compressing  into  tanks  the  oxygen  which 
evaporates  last.  This  is  a  purely  mechanical  process. 

2.  There  are  many  compounds  which,  when  heated  to  tempera- 
tures under  2000°  such  as  we  can  obtain  with  the  aid  of  a  Bunsen 
burner,  a  coal  fire,  or  a  blast-lamp,  give  up  their  oxygen.  Some  of 
them  are  minerals,  but  most  of  them  are  manufactured  articles.  Of 
the  minerals,  pyrolusite  (manganese  dioxide,  MnO2),  employed  by 
Scheele,  is  an  example.  It  usually  contains  the  elements  of  water 
also,  and  hence  moisture  is  evolved  at  the  same  time.  A  substance 
identical  with  the  mineral  hausmannite  (Mn3O4.  For  the  equation 
see  p.  50)  remains.  Amongst  the  artificial  sources  are  mercuric 
oxide,  expensive,  but  historically  interesting  (p.  7);  barium  per- 
oxide, used  in  manufacturing  oxygen  on  a  large  scale  (Brin's  process) ; 
and  potassium  chlorate,  the  most  convenient  for  laboratory  use. 

Erin's  Oxygen  Process  starts  from  barium  oxide  (q.v.).  Barium 
oxide  BaO  closely  resembles  quicklime  CaO,  but  differs  from  this  sub- 
stance in  the  fact  that  when  heated  in  air  to  about  500°,  it  rapidly 
acquires  additional  oxygen  and  gives  barium  peroxide.  When 
barium  peroxide  is  raised  to  1000°,  this  extra  oxygen  is  given  up 
again.  Barium  oxide  contains  one  chemical  unit  weight  each  of  the 
two  constituents  and  takes  up  another  of  oxygen,  so  that  the  equa- 
tion for  the  primary  action  is:  BaO  +  O  ->  Ba02.  The  subsequent 
decomposition  of  the  peroxide,  during  the  stage  in  which  the  oxygen 
is  made,  is  the  exact  opposite:  BaO2  -»  BaO  +  O.*  The  commer- 

*  In  cases  where  an  action  is  reversible,  and  the  direction  depends  on 
conditions  which  may  be  altered,  we  write  both  equations  in  one: 

BaO  +  O  <±  Ba02. 


OXYGEN  47 

cial  advantage  of  the  method  lies  in  the  fact  that  the  barium  oxide 
remaining  after  the  second  stage  can  be  used  over  and  over  again. 
This,  as  will  be  seen,  is  in  reality  a  chemical  method  of  obtaining 
oxygen  from  the  air. 

In  practice  the  barium  oxide  is  maintained  at  a  temperature  of 
700°,  intermediate  between  the  two  just  mentioned,  and  air  is  forced 
under  pressure  into  the  tubes  containing  the  oxide.  A  valve  at  the 
extremity  of  the  tubes  permits  the  escape  of  the  nitrogen.  When 
the  combination  with  oxygen  is  completed,  the  pumping  apparatus 
is  reversed,  and,  a  partial  vacuum  being  created,  the  oxygen  in  com- 
bination is  given  off  without  any  alteration  in  temperature  being 
necessary.  Thus  a  great  waste  of  fuel  is  avoided,  and  the  process 
is  rendered  more  nearly  continuous.  This  method  furnishes  oxygen 
about  96  per  cent  pure,  and  suitable  for  sale  in  compressed  form  in 
iron  cylinders. 

Potassium  Chlorate  (q.v.)  is  a  white  crystalline  substance  used  in 
large  quantities  in  the  manufacture  of  matches  and  fireworks.  When 
heated  in  a  tube  similar  to  that  in  Fig.  5,  it  first  melts  (351°)  and 
then,  on  being  more  strongly  heated,  it  effervesces  and  gives  off  a 
very  large  volume  of  oxygen.  Examination  shows  that  the  whole 
of  the  oxygen  it  contains  can  be  driven  out.  The  white  material 
which  remains  after  the  heating  is  identical  with  the  mineral  sylvite. 
To  the  chemist  it  is  known  as  potassium  chloride,  and,  when  de- 
composed, it  yields  potassium  and  chlorine  in  the  exact  ratio  of  their 
atomic  weights.  Its  formula  is  thus  KC1.  We  may  infer,  therefore, 
that  the  composition  of  the  original  substance  will  be  representable 
by  the  formula  KC1OX,  where  x  is  the  number  of  atomic  weights  of 
oxygen.  Measurement  and  calculation  show  x  =  3.  The  formula 
is  therefore  KC103,  and  the  equation  for  the  decomposition  (see, 
however,  under  Perchlorates) : 

KC1O3  -»  KC1  +  3O. 

A  peculiarity  of  this  action  is  that  admixture  of  manganese  dioxide 
increases  very  markedly  the  speed  with  which  the  decomposition  of 
the  potassium  chlorate  takes  place.  Hence,  in  its  presence,  and  it  is 
generally  mixed  with  the  chlorate  in  laboratory  experiments  (Fig. 
16),  a  sufficient  stream  of  the  gas  is  obtained  at  a  relatively  low 
temperature  (below  200°  see  p.  55). 


48 


COLLEGE    CHEMISTRY 


Physical  Properties  of  Oxygen.  —  Oxygen  resembles  air  in 
being  a  colorless,  tasteless,  and  odorless  gas.  Its  density,  using  the 
physical  standard  of  air  =  1,  is  1.105.  If  hydrogen  is  the  standard, 
oxygen  is  15.900  times  (Morley)  as  heavy.  One  liter  of  oxygen,  at 
0°  and  760  mm.  barometric  pressure,  weighs  1.42900  grams  (Morley). 
The  gas  dissolves  to  some  extent  in  water,  the  solubility  at  0°  being 


FIG.  16 

four  volumes  of  gas  in  one  hundred  volumes  of  water  (at  20°,  3  :  100). 
Liquid  oxygen,  which  was  first  made  by  Wroblewski,  has  a  pale-blue 
color,  and  boils  under  one  atmosphere  at  —  182.5°.  Its  specific 
gravity  at  —  182.5°  is  1.13  (water  =  1):  that  is  to  say,  1  c.c.  weighs 
1.13  g.  By  cooling  with  a  jet  of  liquid  hydrogen,  Dewar  froze  the 
liquid  to  a  snow-like,  pale-bluish  solid.  A  tube  of  liquid  oxygen  is 
very  noticeably  attracted  by  a  magnet. 

Chemical  Properties  of  Oxygen.  —  Sulphur,  when  raised  in 
advance  to  the  temperature  necessary  to  start  the  action,  unites 
vigorously  with  oxygen  (Fig.  17),  giving  out  much  heat  and  pro- 
ducing a  familiar  gas  having  a  pungent  odor  (sulphur  dioxide). 
This  odor  is  frequently  spoken  of  as  the  "  smell  of  sulphur,"  but  in 
reality  sulphur  itself  has  no  odor,  and  neither  has  oxygen.  The  odor 
is  peculiar  to  the  compound  of  the  two.  The  mode  of  experimenta- 
tion can  be  changed  and  the  oxygen  led  into  sulphur  vapor  through 
a  tube.  The  former  then  appears  to  burn  with  a  bright  flame, 
giving  the  same  product  as  before.  For  the  formulae  and  equations 
for  this  and  the  following  actions,  see  next  section. 


OXYGEN 


49 


Warm  phosphorus  combines  with  oxygen  with  even  greater  vigor, 
and  forms  a  white,  powdery,  solid  compound  (phosphoric  anhydride), 
which  absorbs  moisture  from  the  aqueous  vapor 
in  the  air  and  quickly  forms  a  solution  in  this 
water.  In  both  these  cases  the  products  differ 
from  oxygen,  not  only  in  odor  and  in  other 
physical  properties,  but  notably  in  that,  when 
shaken  with  water,  they  dissolve  and  interact 
to  form  acids  (see  below). 

Burning  carbon,  in  the  form  of  charcoal, 
glows  much  more  brightly  in  oxygen  than  in 
ordinary  air.  The  product  is  an  odorless  gas 
(carbon  dioxide).  When  this  gas  is  shaken 
with  "  lime-water, "  a  solution  of  calcium 
hydroxide Ca(OH)2  (q.v.,  and  seep.  71),  a  white 
precipitate  of  calcium  carbonate  CaCO3  is 
formed. 

Finally,  metallic  iron,  which  is  simply  rusted 
by  air  (diluted  oxygen),  burns  in  pure  oxygen 
with  surprising  brilliancy.  Globules  of  a 
molten  product  fall  from  the  iron  and,  when 
they  have  cooled,  are  found  to  consist  of  a 
dark-gray  brittle  material,  which  we  recog- 
nize as  identical  with  blacksmith's  hammer 
scale  and  with  a  well-known  ore  of  iron  (magnetic  oxide  of  iron). 

The  results  of  a  long  series  of  experiments  like  the  above  enable 
us  to  summarize  the  chemical  properties  of  oxygen.  The  gas  unites 
directly  with  nearly  all  the  simple  substances,  and  often,  though  not 
always,  with  the  same  vigor  as  in  the  case  of  these  examples.  In 
the  case  of  one  or  two  elements,  such  as  gold  and  platinum,  the 
compounds  are  obtainable  by  double  decomposition,  and  not 
directly.  With  the  five  members  of  the  helium  group,  of  which  no 
chemical  compounds  are  known,  and  with  fluorine,  oxygen  does 
not  combine. 

Oxygen  can  unite  with  many  of  the  same  elements  when  they  are 
already  in  combination.  Wood,  for  example,  is  composed  of  carbon 
and  hydrogen,  with  a  certain  amount  of  oxygen.  When  previously 
heated,  it  is  decomposed,  and  the  constituents  unite  with  oxygen 
forming  carbon  dioxide  and  water. 


FIG.  17 


50 


COLLEGE    CHEMISTRY 


The  Making  of  Equations  Again.  —  To  learn  the  exact  nature 
of  interactions  like  those  used  as  illustrations  above,  quantitative 
experiments  must  of  course  be  made.  Thus,  for  example,  a  known 
weight  of  sulphur  is  placed  in  a  porcelain  boat  (Fig.  18),  which  has 


I               / 

a 

FIG.  18 


already  been  weighed.  The  U-shaped  tube  to  the  right  contains  a 
solution  of  potassium  hydroxide  which  is  capable  of  absorbing  the 
resulting  gas.  The  oxygen  enters  from  the  left.  When  the  sulphur 
is  heated,  it  burns  in  the  oxygen,  and  the  loss  in  weight  which  the 
boat  undergoes  shows  the  amount  of  sulphur  consumed.  The  gain 
in  weight  of  the  U-tube  shows  the  weight  of  the  compound  produced. 
By  subtracting,  we  get  the  quantity  of  oxygen.  The  proportion  of 
the  constituents  and  the  steps  in  the  calculation  (p.  41)  are  as 
follows: 


PERCENTAGE      AT.  WT.         FACTOR 
Sulphur,          50.05     =     32.06     X     1.561      = 
Oxygen,          49.95     =     16.00      X     3.122     = 


•M.561 

S  X  1.561        S  X  1 
O  X  3.122       O  X  2 


The  formula  of  the  product  is  therefore  S02  and  the  equation  S  +  2O 
->S02. 

Similarly,  phosphoric  anhydride  may  be  shown  to  have  the  for- 
mula P2O5,  carbon  dioxide  CO2,  and  magnetic  oxide  of  iron  Fe3O4. 
To  make  the  equations  representing  their  formation,  the  rules  given 
on  pp.  41-42  should  be  applied. 

To  make  the  equation  for  the  formation  of  oxygen  and  hausman- 
nite  by  heating  manganese  dioxide  (p.  46)  the  same  rules  are  used. 
The  skeleton  equation  is:  Mn02  — >Mn304  +  0.  These  being  the 
formulae  of  the  substances  actually  concerned,  in  balancing  the 
equation  it  is  not  permissible  to  alter  them  except  by  placing 
coefficients  in  front.  The  most  complex  formula  is  Mn304  and,  to 
get  a  sufficient  number  of  atomic  weights  to  produce  it,  at  least 


OXYGEN  51 

3MnO2  is  required.  This  furnishes  3Mn  and  6O,  so  that,  after  using 
40,  there  remains  20 : 

3MnO2  ->  Mn304  +  20. 

Oxides.  —  Substances  containing  one  element  in  combination  with 
oxygen  are  called  oxides,  and  processes  like  those  described  above  are 
called  oxidizing  processes,  or  oxidations.  When  the  same  element 
forms  more  than  one  oxide,  the  names  of  the  oxides  indicate  the 
differing  proportions.  Thus  we  have  barium  oxide  (or  monoxide) 
BaO,  and  barium  peroxide  (or  dioxide)  Ba02,  magnetic  oxide  of 
iron  Fe304,  ferrous  oxide  FeO,  and  ferric  oxide  Fe203.  In  cases  like 
the  last  two  the  terminations  -ous  and  -ic  applied  to  the  metal  corre- 
spond to  the  smaller  and  larger  proportions  of  oxygen,  respectively, 
which  the  metal  is  able  to  hold  in  combination. 

Many  oxides,  like  those  of  iron,  are  almost  indifferent  to  water, 
but  others,  like  those  of  sulphur  and  phosphorus,  interact  with  it 
readily  (see  under  Water).  Some  give  sour  solutions,  containing 
acids  dissolved  in  the  excess  of  water.  Such  solutions  turn  blue 
litmus,  a  vegetable  dye,  red.  Others  give  solutions  with  a  taste  like 
soap  or  borax,  and  here  the  dissolved  substance  is  called  a  base  (q.v.), 
and  turns  litmus  blue.  Thus  sulphur  dioxide  and  phosphoric  anhy- 
dride give  sulphurous  acid  and  phosphoric  acid  respectively: 

S02  +  H20  <=>  H2S03 
P205  +  3H20->2H3P04.* 

If  the  product,  to  whichever  class  it  belongs,  is  not  volatile,  it  may 
be  obtained  by  evaporating  the  excess  of  water.  In  the  case  of 
sulphurous  acid,  the  above  action  is  reversed  by  evaporation  and  the 
sulphur  dioxide  and  water  both  pass  off;  in  that  of  phosphoric  acid, 
the  white  crystalline  acid  is  obtained.  In  consequence  of  their  rela- 
tion to  the  acid,  differing  from  it  in  not  containing  the  elements  of 
water,  these  oxides  are  often  called  anhydrides. 

Combustion.  —  Violent  union  with  oxygen  is  called  in  popular 
language  combustion  or  burning.  Yet  since  oxygen  is  only  one  of 
many  gaseous  substances  known  to  the  chemist,  and  similar  vigorous 
interactions  with  these  gases  are  common,  the  term  has  no  scien- 

*  Here  summation  of  the  formulae  on  the  left  would  give  H6P2O8,  but  in 
such  cases,  unless  there  are  reasons  to  the  contrary,  the  common  factor  is  put 
in  front  and  the  formula  reduced  to  its  lowest  terms. 


52  COLLEGE    CHEMISTRY 

tific  significance.  The  union  of  iron  and  sulphur,  even,  gives  out 
light  and  heat,  and  is  quite  similar  in  the  chemical  point  of  view  to 
combustion. 

In  connection  with  this,  however,  it  may  be  worth  while  to  notice 
the  distinction  between  combustible  and  incombustible  substances. 
Things  which  are  incombustible  may  be  divided  into  two  classes. 
There  are  those  substances  which  already  contain  all  the  oxygen 
which  they  can  hold  in  combination.  Such  are  the  oxides  whose 
formation  we  observed  in  the  experiments  described  above.  In 
everyday  life,  limestone,  sand,  bricks,  and  most  rocks  are  illustra- 
tions. The  other  substances  ordinarily  classed  as  incombustible  are 
those  which  do  not  unite  with  diluted  oxygen,  as  it  is  found  in  the 
air,  with  sufficient  vigor.  The  iron  used  in  the  construction  of 
fireproof  buildings  is  the  commonest  example  of  this  class. 

Oxidation.  —  The  rusting  of  metals  differs  from  combustion 
mainly  in  speed.  Often  the  products  are  identical  in  composition 
and  properties  with  the  oxides  formed  by  combustion.  In  the  case 
of  iron,  burning  gives  us  the  magnetic  oxide  (Fe304),  while  rusting 
in  cold,  moist  air  yields  a  hydrated  ferric  oxide  (Fe203  +  Aq*). 
The  products  differ  in  composition,  but  are  closely  related. 

This  process  of  slow  oxidation,  although  less  conspicuous  than 
combustion,  is  really  of  greater  interest.  Thus  the  decay  of  wood  is 
simply  a  process  of  oxidation  whereby  the  same  products  are  formed 
as  by  the  more  rapid  ordinary  combustion.  Large  volumes  of  pure 
water  are  mixed  with  sewage,  the  object  being,  not  simply  to  dilute 
the  latter  but  to  introduce  water  containing  oxygen  in  solution. 
This  has  an  oxidizing  power  like  that  of  oxygen  gas,  and,  through 
the  agency  of  bacteria,  quickly  renders  dissolved  organic  matters 
innocuous  by  converting  them  for  the  most  part  into  carbon  dioxide 
and  water.  In  our  own  bodies  we  have  likewise  a  familiar  illustra- 
tion of  slow  oxidation.  Avoiding  details,  it  is  sufficient  to  say  that 
the  oxygen  from  the  air  taken  into  the  lungs  combines  with  the 
haemoglobin  in  the  red  blood-corpuscles.  In  this  form  of  loose 
combination,  it  is  carried  by  the  blood  throughout  our  tissues  and 

*  The  formula  H2O  may  not  be  used  excepting  to  indicate  a  definite  pro- 
portion of  the  elements  of  water  (18  parts).  Where  the  proportion  varies 
according  to  circumstances,  as  here  and  in  the  case  of  solutions,  the  contrac- 
tion Aq  is  employed. 


OXYGEN  53 

is  there  used  for  oxidizing  waste  materials.  The  carbon  dioxide  is 
carried  back  to  the  lungs  by  the  blood,  and  finally  reaches  the  air 
during  exhalation.  To  supply  the  place  of  the  material  thus 
removed,  we  are  under  the  continual  necessity  of  building  new  tissue 
from  the  food  which  we  eat.  If  we  cease  to  eat,  we  become  lighter 
and  weaker,  showing  that  a  real  portion  of  our  structure  is  gradually 
being  consumed  by  oxidation. 

The  opposite  of  oxidation,  the  removal  of  oxygen,  is  spoken  of  in 
chemistry  as  reduction. 

Means  of  Altering  the  Speed  of  a  Given  Chemical  Action : 
By  Change  of  Temperature.  —  That  the  same  change  may  pro- 
ceed with  very  different  speeds  according  to  conditions  is  a  familiar 
fact.  For  example,  raising  the  temperature  increases  the  rapidity  of  all 
chemical  interactions.  Thus,  cold  iron  combines  with  oxygen  very 
slowly,  giving  rust,  while  white-hot  iron  sheds  quantities  of  scales  of 
an  oxide,  formed  in  the  few  moments  that  it  is  under  the  blacksmith's 
hammer.  White-hot  coal  unites  with  oxygen  in  the  air  to  form 
carbon  dioxide  and  seems  to  disappear  before  our  eyes,  while  in  the 
cellar,  even  in  warm  weather,  we  observe  no  appreciable  diminution 
in  its  amount.  No  temperature  can  be  found,  however,  at  which 
the  interaction  definitely  begins.  We  believe  that  every  such 
change  proceeds  with  some  speed  at  every  temperature. 

If,  on  bringing  two  materials  together,  the  chemist  observes  no 
marks  of  chemical  action,  he  immediately  begins  cautiously  to  heat 
the  mixture.  This  appeal  to  the  magnifying  effect  of  a  rise  in  tem- 
perature is  always  made  as  a  matter  of  course. 

The  common  expressions  used  in  chemistry  in  describing  temper- 
atures, along  with  the  corresponding  readings  of  the  thermometer, 
are  as  follows: 

Incipient  red  heat,  about  525°.     Yellow  heat,  about  1100°. 

Dark  red  heat,  "     700°.     Beginning  white  heat,     "      1300°. 

Bright  red  heat,  "    950°.     White  heat,  "      1500°. 

Rapid  Self -sustaining  Chemical  Action  and  Means  of  Ini- 
tiating it.  —  When  a  piece  of  wood  is  set  on  fire  at  one  end,  the  heat 
produced  by  the  action  itself  raises  the  temperature  of  neighboring 
portions  until  their  speed  of  union  becomes  equal  to  that  of  the  part 
originally  lighted.  In  this  way  the  whole  becomes  finally  inflamed. 


54  COLLEGE   CHEMISTRY 

When  we  blow  the  blaze  out,  the  great  excess  of  cold  air  suddenly 
lowers  the  temperature  of  the  wood,  and  of  the  gas  rising  from  it,  and 
rapid  union  ceases.  Whether  a  given  set  of  materials  can  maintain 
itself  at  a  temperature  proper  to  violent  interaction  will  depend  on 
the  amount  of  heat  developed  by  the  action  itself  on  the  one  hand, 
and  the  losses  of  heat  by  conduction  and  radiation  on  the  other.  If 
the  latter  are  great,  the  former  must  be  greater.  Thus  the  union  of 
iron  and  oxygen  per  se  gives  heat  enough  to  warm  the  materials  to 
the  burning  temperature  and  leaves  much  over  for  radiation.  But 
iron  in  air,  which  is  four-fifths  nitrogen,  can  receive  the  oxygen  only 
one-fifth  as  fast  at  the  start,  and  even  more  slowly  as,  later,  the 
nitrogen  accumulates  round  it.  And  besides,  all  the  nitrogen  has 
to  be  heated  to,  perhaps,  2000°.  The  task  is  too  great.  The  union 
is  impeded  and  the  iron  is  not  oxidized  fast  enough  to  generate  the 
heat  required  to  maintain  everything  at  this  high  temperature. 
Poor  conductors  of  heat,  like  wood  and  candles,  fare  better.  Pow- 
dered iron,  with  its  particles  presenting  large  surface  to  the  air 
relatively  to  the  weight  of  material  in  each  particle  to  be  heated, 
burns  well. 

The  initial  supply  of  heat  required  to  start  a  violent  chemical 
action  must  not  be  confused  with  the  heat  subsequently  developed 
as  the  action  proceeds.  The  preliminary  supply  varies  with  circum- 
stances, and  may  be  made  as  small  as  we  choose  by  limiting  the  area 
first  heated  and  using  ordinary  precautions  against  radiation  and 
convection.  The  heat  produced  by  the  interaction  itself,  however,  is 
fixed  in  amount,  and  depends  only  on  the  materials  and  their  quantity. 

Heating  is  not  the  only  means  used  to  give  the  initial  acceleration 
to  a  self-sustaining  chemical  change.  Thus  in  striking  a  match  a 
rather  violent  vibration  is  employed  to  hasten  the  torpid  action  in 
a  small  part  of  the  material,  and  the  heat  produced  by  the  resulting 
action  quickly  ignites  the  whole.  The  same  explanation  accounts 
for  the  explosion  of  gun-cotton  by  a  percussion  fuse. 

Other  Means  of  Altering  the  Speed  of  a  Given  Chemical 
Change :  By  Catalysis.  —  When,  without  any  change  in  tempera- 
ture, an  extra  substance  increases  the  speed  of  a  chemical  change, 
seemingly  by  its  mere  presence,  without  itself  suffering  any  perma- 
nent change,  we  call  this  catalytic  (Gk.  Kara ,  down;  Arms,  the  act  of 
loosing)  or  contact  action.  The  word  was  originally  used  for  cases  of 


OXYGEN  55 

decomposition.  The  foreign  body  is  called  the  catalytic  or  contact 
agent,  and  the  process  catalysis.  The  effect  of  manganese  dioxide 
on  the  decomposition  of  potassium  chlorate  (p.  47)  is  of  this  nature. 
When  some  of  the  chlorate  is  placed  in  a  test-tube,  provided  with  a 
two-hole  stopper  and  delivery  tube,  and  is  melted  carefully  so  as  to 
avoid  superheating,  scarcely  any  evolution  of  oxygen  can  be  per- 
ceived at  this  temperature  (351°).  If  now  a  pinch  of  pulverized 
manganese  dioxide,  hitherto  held  in  one  of  the  holes  of  the  stopper 
by  means  of  a  small  plug  of  glass  rodding,  be  dropped  into  the 
molten  mass  by  forcing  a  longer  rod  into  the  hole,  the  oxygen  is 
given  off  in  torrents  in  consequence  of  the  enormous  acceleration  of 
the  decomposition.  Yet  the  manganese  dioxide  may  be  recovered 
unchanged  from  the  residue.  Manganese  dioxide,  of  course,  will 
itself  give  oxygen  (p.  46),  but  the  decomposition  is  hardly  noticeable 
at  400°. 

[/  Thermochemistry.  —  As  we  have  seen,  heat  is  liberated  in  con- 
nection with  many  chemical  changes  (pp.  6,  13,  49).  Such  changes 
are  called  exothermal.  In  other  chemical  changes  (pp.  8, 13,  46)  heat 
is  continuously  absorbed.  These  are  called  endothermal  changes. 
Since  the  activities,  or  affinities  of  two  substances  (say,  two  metals) 
may  often  be  measured  (p.  19)  by  observing  the  amounts  of  heat 
liberated  when  each  combines  with  a  third  substance  (say,  oxygen), 
it  will  be  instructive  now  to  consider  some  of  the  elementary  facts 
of  thermochemistry. 

The  chemical  interactions  to  be  studied  thermally  arc  arranged  so 
that  they  may  be  carried  out  in  some  small  vessel  which  can  be 
placed  inside  another  containing  water.  The  whole  apparatus  is 
called  a  calorimeter.  The  heat  developed  raises  the  temperature  of 
this  water.  Where  gases  like  oxygen  are  concerned,  a  closed  bulb 
of  platinum  forms  the  inner  vessel.  The  quantity  of  heat  capable 
of  raising  one  gram  of  water  one  degree  in  temperature,  between 
0°  and  100°  Centigrade,  is  called  a  calorie.  So  that  250  grams  of 
water  raised  1°  would  represent  250  calories,  and  20  grams  of  water 
raised  5°  would  represent  100  calories. 

While  in  physics  the  unit  of  quantity  is  the  gram,  in  chemistry 
the  unit  which  we  select  is  naturally  that  represented  by  the  for- 
mula of  the  substance.  Thus,  the  heat  of  combustion  of  carbon 
means  the  heat  produced  by  combining  twelve  grams  of  carbon  with 


56  COLLEGE   CHEMISTRY 

thirty-two  grams  of  oxygen,  and  is  sufficient  to  raise  nearly  100,000 
grams  of  water  one  degree.  This  is  expressed  as  follows: 

C  +  2O  -»  C02  +  100,000  cal. 

In  other  words,  the  combustion  of  less  than  half  an  ounce  of  carbon 
will  raise  one  kilogram  (over  two  pounds)  of  water  from  0°  to  the 
boiling-point. 

It  is  always  found  that  the  same  quantities  of  any  given  chemical 
substances  sustaining  the  same  chemical  change  under  the  same  con- 
ditions produce  or  absorb,  according  as  the  action  is  exothermal  or 
endothermal,  amounts  of  heat  which  are  equal. 

The  rate  at  which  a  given  chemical  action  is  allowed  to  take  place 
has  no  influence  on  the  total  amount  of  heat  consumed  or  produced. 
It  may  not  at  first  sight  appear  obvious  that  rusting  evolves  heat, 
but  a  delicate  thermometer  will  show  that  a  heap  of  rusting  nails  is 
somewhat  higher  in  temperature  than  surrounding  bodies.  Poor 
conductors,  like  oily  rags  and  ill-dried  hay,  show  a  tendency  to  spon- 
taneous combustion  owing  to  accumulation  of  the  slowly  developing 
heat  of  oxidation.  The  warmth  of  our  own  bodies  is  in  part  due  to 
the  same  cause. 

It  should  be  noted  that  production  or  absorption  of  heat  is  not, 
in  itself,  an  evidence  of  chemical  action.  Physical  changes  are  all 
likewise  accompanied  by  the  same  phenomena.  Thus,  the  evapora- 
tion of  water  absorbs  heat,  and  condensation  of  a  vapor  and  the 
crystallization  of  a  supercooled  liquid  liberate  heat. 

Exercises.  —  1.  Enumerate  other  instances,  already  encount- 
ered, of  the  use  of  the  terminations  ous  and  ic  to  distinguish 
different  degrees  of  oxidation.  For  what  other  purposes  have 
the  same  terminations  been  used? 

2.  What  difference  in  composition  between  potassium  chloride  and 
chlorate  are  the  terminations  designed  to  indicate?     Applying  the 
same  idea,  how  would  ferrous  sulphate  (q.v.)  differ  from  ferrous 
sulphide,  and  cupric  sulphate  from  cupric  sulphide? 

3.  Define  and  illustrate:  density  of  a  gas  (p.  48,  and  see  p.  61), 
specific  gravity  of  a  solid  or  liquid  (pp.  2,  23). 

4.  Enumerate  the  classes  of  facts  given  under  the  heads  of,  Phys- 
ical Properties,  and  Chemical  Properties  of  oxygen,  respectively 
(see  p.  75). 


OXYGEN  57 

5.  Construct  the  equations  for  the  combustion  of  phosphorus, 
carbon,  and  iron  in  oxygen  (pp.  49,  50). 

6.  When  1  g.  of  sodium  burns  in  oxygen,  it  produces  1.7  g.  of  the 
oxide.     What  is  the  formula  of  the  latter,  and  the  equation  (p.  41)  ? 

7.  Which  are  the  components  (p.  22)  of  the  liquid  made  by  treat- 
ing phosphoric  anhydride  with  water?    Which  are  the  constituents 
(p.  51)  of  phosphoric  acid? 

8.  How  should  you  show  that,  in  the  making  of  oxygen  from  a 
mixture  of  potassium  chlorate  and  manganese  dioxide,  the  latter 
remains  unchanged?     Which  properties  (p.  23)  are  you  employing 
for  this  purpose? 

9.  The  substances,   like  phosphorus  and  sulphur,   which  burn 
rapidly  in  ordinary  oxygen,  combine  very,  very  slowly  with  oxygen 
which  has  been  freed  from  moisture  by  careful  drying.     How  is  this 
effect  of  water  to  be  classified? 

10.  Discuss  the  union  of  iron  and  sulphur  (p.  6)  and  the  decom- 
position of  mercuric  oxide  (p.  7)  in  their  relation  to  the  explana- 
tions on  pp.  53-54. 

11.  How  many  calories  are  required  to  raise  500  g.  of  a  substance 
of  specific  heat  0.5  from  15°  to  37°  (p.  55)? 

12.  The  combustion  of  1  g.  of  sulphur  to  sulphur  dioxide  develops 
2220  calories.     What  is  the  heat  of  combustion  of  sulphur  (p.  55)  ? 


CHAPTER  VI 


THE    MEASUREMENT    OF   QUANTITY   IN    GASES 

A  SPECIMEN  of  a  gas,  like  a  specimen  of  a  solid  or  a  liquid,  may  be 
weighed,  but  it  is  usually  easier  to  determine  the  quantity  of  the 
gas  by  (1)  measuring  its  volume,  and  at  the  same  time  (2)  noting  its 
temperature  on  a  thermometer  suspended  in  it  or  close  to  it,  and  (3) 
ascertaining  the  pressure  which  it  exercises. 

The  Measurement  of  the  Pressure  of  a  Gas.  —  In  almost  all 
cases  the  easiest  way  to  take  account  of  the  pressure  of  a  gas  is  to 
place  it  in  an  apparatus  so  constructed  that  one 
boundary  of  the  volume  is  a  liquid.  The  appar- 
atus is  then  so  adjusted  that  the  surface  of  the 
liquid  in  contact  with  the  gas  is  at  the  same  level 
as  the  free  surface  of  the  liquid  which  is  exposed 
to  the  atmosphere.  The  equality  in  the  levels 
of  the  liquids  is  then  a  guarantee  that  the 
specimen  of  gas  and  the  atmosphere  are  exercis- 
ing equal  pressures  on  the  liquid.  At  this 
stage  the  volume  of  the  gas  is  measured,  and 
simultaneously  the  pressure  of  the  atmosphere 
and,  therefore,  of  the  gas,  is  ascertained  by 
reading  the  barometer. 

The  barometer  (Fig.  19)  consists  of  a  bent 
tube  containing  mercury.  The  short  limb 
(to  the  left)  is  open  and  the  pressure  of  the 
atmosphere  is  exercised  on  the  surface  of  the 
mercury  there.  The  longer  limb  (to  the  right) 
is  closed  at  the  top  and  in  it  there  is  no  gas 
above  the  mercury.  When  the  tube  is  inclined, 
Fjo  19  the  surface  of  the  mercury  in  the  longer  limb 

endeavors  to  retain  the  same  vertical   height 
above  the  lower  surface  and  consequently  rises,  and,  with  sufficient 

68 


MEASUREMENT  OF   QUANTITY  IN  GASES  59 

inclination,  will  reach  entirely  to  the  top  of  the  tube.  The 
downward  pressure  of  the  mercury  on  the  right,  above  the  dotted 
line,  is  exactly  equal  to  the  pressure  of  the  atmosphere  on  the  free 
surface  of  the  mercury  at  the  same  level.  The  amount  of  the  latter 
pressure  is  proportional  to  the  length  of  the  column  of  mercury  above 
the  dotted  line.  Hence,  reading  the  height  at  which  the  mercury 
stands  above  the  free  surface  gives  us  a  measure  of  the  pressure  of  the 
atmosphere  and  of  any  specimen  of  gas  which  is  at  the  same  pressure. 

This  is  called  the  uncorrected  reading.  It  is  immediately  reduced 
to  the  reading  which  would  have  been  made  if  the  barometer  and  its 
mercury  had  been  at  0°  (corrected  reading),  by  noting  the  temperature 
on  the  adjacent  thermometer  and  subtracting  from  the  uncorrected 
reading  the  necessary  correction  (Table  of  corrections,  C,  Fig.  19). 

For  example:  the  volume  of  gas,  after  adjustment  to  atmospheric 
pressure,  is  200  c.c.  and  its  temperature  17°.  The  uncorrected  baro- 
metric reading  is  744  mm.  with  the  barometer  (perhaps  in  a  different 
room  from  the  gas)  at  15°.  The  correction  is  -  2.0  mm.  The  cor- 
rected reading  is  therefore  742  mm. 

Finally,  since  the  atmospheric  pressure  varies  from  day  to  day, 
the  volume  at  the  observed  pressure  is  corrected  to  that  which  the 
same  quantity  of  gas  would  have  occupied  at  the  standard  pressure 
of  760  mm.  of  mercury.  By  careful  measurements,  Boyle  (1660) 
found  that  the  volumes  occupied  by  the  same  sample  of  any  gas  are 
inversely  proportional  to  the  pressures  at  each  volume. 

The  illustration  just  given  will  show  how  this  additional  correction 
is  applied.  There  were  200  c.c.  of  the  gas  at  17°  and  742  mm. 
(corr.).  The  question  is:  What  would  be  the  volume  of  this  amount 
of  gas  at  760  mm.?  At  this  new  pressure  (760  mm.)  which  is  greater 
than  the  old  pressure  (742  mm.),  the  volume  will  become  less. 
Hence  we  change  the  volume  in  the  proportion  of  these  pressures, 
placing  the  smaller  number  in  the  numerator,  so  as  to  get  a  smaller 
volume  as  the  answer:  200  X  £$$  =  volume  at  760  mm.,=  195.3  c.c. 
If  we  wished  to  convert  100  c.c.  at  775  mm.  to  760  mm.,  we  should 
reason  that  the  new  pressure  was  smaller,  and  the  volume  would 
become  greater,  and  should  therefore  place  the  larger  number  (775) 
in  the  numerator  so  as  to  get  a  larger  volume  for  the  answer. 

The  Correction  of  the  Volume  of  a  Gas  for  Temperature.  — 
The  same  sample  of  gas  will  occupy,  when  heated,  a  larger  volume, 
and  when  cold,  a  smaller  volume  than  before.  The  change  in  volume 


60  COLLEGE    CHEMISTRY 

for  each  degree  Centigrade  is  -%\^  of  the  volume  of  the  same  sample 
at  0°.  To  simplify  the  calculation  we  begin  by  converting  the  tem- 
perature to  the  absolute  scale  by  adding  273°  to  each  temperature: 
The  volumes  assumed  by  a  sample  of  gas  at  different  temperatures,  the 
pressure  remaining  constant,  are  in  the  same  proportion  as  the  correspond- 
ing absolute  temperatures  (Charles,  1787).  If  the  volume  remains 
constant,  then  the  pressure  changes  in  the  same  proportion. 

In  the  illustration  used  above,  there  were  200  c.c.  of  gas  at  17°, 
and  it  is  required  to  know  the  volume  at  0°.  We  add  273  algebra- 
ically to  each  temperature,  and  the  question  becomes:  There  are 
200  c.c.  of  gas  at  290°  Abs.,  what  will  be  its  volume  at  273°  Abs.? 
The  volume  changes  in  the  direct  ratio  of  the  temperatures.  The 
new  temperature  is  lower  than  the  old,  and  the  new  volume  will 
therefore  be  smaller  than  the  old.  Then  200  X  f-£J  =  volume  at 
0°  (273°  Abs.)  =  188.27  c.c. 

The  above  laws  are  usually  applied  to  any  example  simultaneously. 
Thus,  200  c.c.  of  gas  at  742  mm.  pressure  (corr.)  and  17°  become 
200  X  £ff  X  f|f  =  183.8  c.c.  at  0°  and  760  mm. 

Mixed  Gases :  Aqueous  Tension.  —  Two  gases  at  the  same 
temperature,  provided  they  do  not  interact  chemically,  do  not  inter- 
fere with  each  other's  pressures  when  mixed.  Thus,  if  they  are  forced 
into  the  same  volume,  the  pressure  of  the  mixture  is  equal  to  the 
sum  of  those  of  the  components  (Dalton's  law,  1807).  The  gases 
are  therefore  still  thought  of  individually,  and  the  share  which  each 
gas  has  in  the  total  pressure  is  called  its  partial  pressure.  This,  like 
any  other  gaseous  pressure,  is  proportional  to  the  concentration  of 
the  particular  gas  in  the  mixture. 

For  example,  a  gas  measured  over  water  contains  water  vapor. 
The  partial  pressure  of  this,  called  the  aqueous  tension  (q.v.),  which 
is  definite  for  each  temperature,  must  be  subtracted  from  the  total 
pressure.  The  remainder  is  the  partial  pressure  of  the  gas  being 
measured,  and  this  remainder  is  used  as  the  pressure  of  this  gas  in 
any  calculation.  Thus,  in  a  gas  measured  over  water  at  22°,  the 
total  pressure  includes  19.7  mm.  (the  aqueous  tension  at  22°)  pressure 
of  water  vapor.  Hence  150  c.c.  of  gas  over  water  at  22°  and  750 
mm.  is  the  same  in  amount  as  150  c.c.  of  the  same  gas  in  dry  condi- 
tion at  22°  and  730.3  mm.  (there  being  simply  150  c.c.  of  water 
vapor  at  19.7  mm.  mixed  with  it).  To  obtain  the  volume  of  dry  gas 
at  0°  and  760  mm.  we  have  the  expression  150  x  f  H  x  *f&3 . 


MEASUREMENT   OF   QUANTITY  IN   GASES  61 

Densities  of  Gases.*  —  The  density  of  a  gas  is  the  mass  of  1  c.c. 
of   the  gas  at  0°  and  760  mm.  pressure.     Sometimes  the  weight 
of  one  liter  (1000  c.c.)   is  called  the  density.     Often  the  relative 
weight   of   the  gas,  the  weight  of  an  equal 
volume  of  air,  or  oxygen,  or  hydrogen  being 
taken  as  unity,  receives  the  same  name. 

The  most  direct  method  of  measuring  the 
density  of  a  gas  is  to  employ  a  light  flask 
provided  with  a  rubber  stopper  and  stopcock 
(Fig.  20).  By  means  of  an  air-pump  the 
contents  of  the  flask  are  removed,  and  it  is 
weighed.  This  gives  the  weight  of  the  empty 
vessel.  The  gas,  whose  density  is  to  be  ascer- 
tained, is  then  admitted,  and  care  is  taken 
that  it  finally  fills  the  flask  at  the  pressure  of 
the  atmosphere.  The  flask  is  closed  and 
weighed  again.  The  increase  represents  the 
weight  of  the  gas.  At  the  same  time  the 
temperature  and  barometric  pressure  are 
read.  The  volume  is  determined  by  dis-  FIQ  2Q 

placing  the  gas  once  more  from  the  flask, 

filling  with  water,  and  weighing  again.  The  difference  in  weight 
between  the  empty  flask  and  the  flask  full  of  water,  in  grams, 
represents  the  volume  of  the  content  of  the  flask  in  cubic  centi- 
meters. This  volume  is  reduced  to  0°  and  760  mm.  by  the  rules 
discussed  above,  and  we  have  then  a  volume  of  the  gas  and  the 
corresponding  weight. 

To  illustrate,  let  us  suppose  that  the  volume  of  the  flask  is  200  c.c. 
and  that  it  is  filled  with  oxygen  at  17°  and  742  mm.  The  weight 
of  the  gas  is  found  to  be  0.26  g.  We  ascertained  (p.  60)  by  calcula- 
tion that  at  0°  and  760  mm.  this  volume  would  be  183.8  c.c.  The 
weight  of  a  liter  is  given  by  the  proportion  183.8  :  0.26  :  :  1000  :  x. 
Here  x  =  1.415  g.  When  the  operation  is  performed  carefully,  and 
the  weighing  carried  to  the  nearest  milligram  instead  of  the  nearest 
centigram,  a  result  more  nearly  approaching  the  exact  one  (1.429) 
may  easily  be  reached. 

To  get  the  density  of  oxygen  referred  to  hydrogen  as  unity,  we 

*  The  subjects  of  this  section  are  not  actually  used  until  Chapter  xii  (on 
Molar  Weights)  is  reached. 


62  COLLEGE  CHEMISTRY 

must  divide  the  answer  by  the  weight  of  a  liter  of  hydrogen  (0.08987 
g.).  In  the  above  case  the  quotient  is  15.74.  The  accepted  value 
is  15.90.  The  density  referred  to  air  as  unity  is  similarly  obtained 
by  dividing  by  1.293,  the  weight  of  a  liter  of  air  at  0°  and  760  mm. 
pressure. 

By  using  a  modification  of  the  flask  just  described,  it  is  possible  to 
ascertain  the  weights  of  known  volumes  of  the  vapors  of  liquids  and 
solids.  A  temperature  sufficiently  high  to  vaporize  the  substance 
must  be  employed.  The  volume  is  reduced  by  rule  to  0°  and  760  mm. 
and  the  density  (in  this  case  known  as  the  vapor  density)  is  calculated 
as  before.  The  reduction  to  0°  and  760  mm.  pressure  by  rule  gives, 
of  course,  a  fictitious  result.  The  vapor  would  condense  to  the  liquid 
form  before  0°  was  reached,  if  the  cooling  were  actually  carried  out. 
But  the  value  for  the  density  as  it  would  be  at  0°  and  760  mm.  has  to 
be  calculated  to  facilitate  comparison  with  the  corresponding  values 
for  other  substances.  The  results  have  no  physical  significance,  but 
are  highly  important  to  the  chemist. 

Exercises.  —  The  foregoing  cannot  be  understood  unless  some 
problems  involving  the  laws  of  gases  are  actually  worked. 

1.  Reduce  189  c.c.  of  gas  at  15°  and  750  mm.  to  0°  and  760  mm. 

2.  Reduce  110  c.c.  of  gas  at  -  5°  and  741  mm.  to  0°  and  760  mm. 

3.  Convert  500  c.c.  of  gas  at  25°  and  700  mm.  to  18°  and  745  mm. 

4.  Reduce  250  c.c.  of  gas  (standing  over  water)  at  22°  and  755 
mm.  to  the  dry  condition  and  to  0°  and  760  mm. 

5.  The  density  of  a  substance  referred  to  air  is  3.2.     What  is  the 
density  referred  to  hydrogen?    What  will  be  the  volume  occupied 
by  10  g.  of  the  substance  at  20°  and  752  mm.  ? 


CHAPTER  VII 
HYDROGEN 

IN  each  chapter  dealing  with  the  chemistry  of  some  substance  the 
same  topics  are  always  discussed  in  the  same  order,  namely,  History, 
Occurrence,  Preparation,  Physical  properties,  and  Chemical  proper- 
ties. Additional  sections  dealing  with  special  topics  are  inserted 
when  necessary. 

The  independent  nature  of  hydrogen  was  first  established  by 
Cavendish  in  1766.  Somewhat  later  (1781),  he  showed  that  hydro- 
gen when  it  burned  gave  water  vapor.  Taken  in  conjunction  with 
Lavoisier's  proof  that  oxygen  was  the  active  substance  in  the  air 
(1777),  this  fact  showed  that  water  was  a  compound  of  hydrogen 
(Gk.  v'Scop,  water;  ywvav,  to  produce)  and  oxygen. 

Occurrence.  —  The  free  element  is  found,  mixed  with  varying 
proportions  of  other  gases,  in  exhalations  from  volcanoes,  in  pockets 
found  in  certain  layers  of  the  rock-salt  deposits,  and  in  some  mete- 
orites. The  air  contains  not  more  than  one  part  in  30,000.  Its  lines 
are  very  prominent  in  the  spectra  of  the  sun  and  of  most  stars. 

In  combination,  it  constitutes  about  1 1  per  cent  of  water.  It  is  an 
essential  constituent  of  all  acids.  It  is  contained  also,  in  combina- 
tion with  carbon,  in  the  components  of  natural  gas,  petroleum,  and 
all  animal  and  vegetable  bodies. 

Acids.  —  In  making  hydrogen,  the  acids  are  used  almost  exclu- 
sively. The  common  acids  are  hydrochloric  acid  (HC1,  Aq),  and 
sulphuric  acid  (H2S04,  Aq).  The  usual  forms  are  mixtures  con- 
taining water,  the  variable  amount  of  the  latter  being  indicated  by 
the  symbol  Aq.  The  former  is  a  solution  of  a  gas,  hydrogen  chloride. 
The  "  pure  concentrated  "  hydrochloric  acid  used  in  laboratories 
contains  nearly  as  much  of  the  gas  (39  per  cent  by  weight)  as  the 
water  can  dissolve.  The  "commercial"  acid  contains  impurities 
and  is  also  less  concentrated.  The  "  concentrated  "  sulphuric  acid 
is  an  oily  liquid  containing  practically  no  water. 

63 


64 


COLLEGE   CHEMISTRY 


cial  "  sulphuric  acid  contains  6  to  7  per  cent  of  water,  besides  impuri- 
ties. Acetic  acid  (HC2H302,  Aq)  is  a  solution  of  a  liquid  in  water. 
All  the  "  dilute  "  acids  contain  70  to  80  per  cent  of  water.  The 
water,  as  a  rule,  takes  no  part  in  the  chemical  changes  in  which  the 
acids  are  concerned,  and  is  therefore  omitted  from  the  equations. 

The  name  "  acid  "  is  restricted  to  one  class  of  substances  having 
certain  definite  characteristics.  Hydrogen  is  the  one  essential  con- 
stituent of  all  acids.  Their  aqueous  solutions  have  a  sour  taste  and 
change  the  color  of  litmus  from  blue  to  red.  When  free  from  water 
they  do  not  conduct  electricity.  When  dissolved  in  water  they  con- 
duct, and  are  decomposed  by  the  electric  current,  and  their  hydrogen 

(or  one  unit  weight  of  it  in  the  case 
of  acetic  acid)  is  displaced  by  certain 
metals. 

In  describing  the  chemical  be- 
havior of  acids,  we  speak  of  the 
hydrogen  as  the  positive  radical, 
because  in  electrolysis  it  is  attracted 
to  the  negative  pole,  and  of  the 
material  combined  with  the  hydrogen 
as  the  negative  radical  because  it  is 
attracted  to  the  positive  pole.  Thus 
the  negative  radicals  in  the  above 
acids  are  Cl,  SO4,  and  C2H3O2,  respec- 
tively. The  first  (Cl)  is  a  simple 
radical,  the  others  compound.  In 
many  interactions  the  compound 
radicals  move  as  units  from  one  state 
of  combination  to  another. 


Preparation  of  Hydrogen  by 
Electrolysis.  —  If  we  dissolve  any 
acid  in  water,  and  immerse  the  wires 
from  a  battery  in  the  solution, 
bubbles  of  hydrogen  begin  to  appear 
on  the  negative  wire  (the  cathode) 

and  rise   to   the  surface.     All   the   other  constituents,  whatever 

'  they  may  be,  are  attracted  to  the  positive  wire  (the  anode)  and 

are  set  free  in  some  form  at  its  surface.     An  apparatus  devised  by 


FlG  21 


HYDROGEN 


65 


Hofmann  (Fig.  21)  enables  us  to  secure  the  hydrogen,  which  ascends 
on  the  left  and  accumulates  at  the  top  of  the  tube,  displacing  the 
solution.  The  other  products,  if  gaseous,  occupy  a  separate  tube 
on  the  right  side.  In  a  typical  case  the  production  of  hydrogen 
ceases  when  the  acid  is  all  decomposed.  The  water  alone  is  an 
almost  complete  nonconductor,  so  that  the  flow  of  the  electricity 
practically  ceases  at  the  same  time.  Thus  when  hydrochloric  acid 
is  used:  HC1  — >  H  (neg.  wire)  +  Cl  (pos.  wire),  and  the  chlorine,  a 
soluble  gas,  remains  dissolved  in  the  water  near  one  pole  (see 
Electrolysis) . 

Preparation  of  Hydrogen  by  Displacement  from  Diluted 
Acids.  —  Certain  metals,  like  zinc, 
iron,  and  aluminium,  when  placed 
in  a  dilute  acid,  combine  with  the 
radical  of  the  acid  and  so  liberate 
the  hydrogen.  The  acids  must  be 
diluted  with  water  before  rapid  action 
occurs.  The  hydrogen  escapes  in 
bubbles,  and  evaporation  of  the 
remaining  liquid  gives  in  dry  form 
the  compound  of  the  metal  with  the 
other  constituents  of  the  acid.  Thus, 
with  zinc  and  sulphuric  .acid,  zinc 
sulphate  is  produced: 

Zn  +  H2S04  ->  2H  +  ZnSO4; 

and    with    tin    or    aluminium    and 
hydrochloric   acid  we  get  stannous 
chloride  or  aluminium  chloride: 
Sn  +  2HC1  ->  2H  +  SnCl2, 

(Tin)  (Stannous  chloride) 

Al  +  3HC1  -»  3H  +  A1C13. 

The  water  undergoes  no  change  dur- 
ing the  action,  although  its  presence 
is  essential.  It  is  simply  a  part  of 

,,  t  j  ,  FIG.  22. 

the  apparatus.      Any  acid  may  be 

used,  although  with  many  the  action  goes  on  very  slowly.  In  all 
cases  the  plan  of  the  action  is  the  same :  the  metal  is  said  to  displace 
the  hydrogen  (see  below). 


66  COLLEGE   CHEMISTRY 

With  a  Kipp's  apparatus  (Fig.  22)  the  gas  may  be  made  on  a 
large  scale  and  its  delivery  can  be  regulated.  When  the  stream  of 
gas  is  shut  off  by  the  stopcock,  the  pressure  of  the  gas,  as  it  con- 
tinues to  be  generated,  drives  the  acid  away  from  the  metal  and  up 
into  the  globe  above,  so  that  the  action  ceases.  Yet  the  action  is 
ready  to  begin  again  the  moment  any  portion  of  the  stored  gas  is 
drawn  off  for  use. 

A  rather  sharp  line  can  be  drawn  between  those  metals  which  dis- 
place hydrogen  from  dilute  acids  and  those  which,  like  mercury, 
silver,  and  gold,  do  not  (see  Electromotive  series,  also  footnote  to 
p.  24).  Contact  of  the  zinc  or  iron  with  an  inactive  metal,  like 
platinum,  forms  an  electrical  couple  and  hastens  the  interaction. 

Preparation  of  Hydrogen  from  Water.  —  When  sodium,  one 
of  the  constituents  of  common  salt,  is  thrown  upon  water,  the  metal 
gradually  disappears  and  hydrogen  is  liberated  with  violent  effer- 
vescence. Every  one  of  the  metals  which  act  on  dilute  acids  will 
displace  hydrogen  from  water,  and  no  others  will  do  so.  The  more 
active  metals,  like  potassium  and  sodium,  which  would  act  with 
uncontrollable  vigor  on  dilute  acids,  displace  the  hydrogen  rapidly 
from  cold  water.  Magnesium  and  zinc  show  obvious  action  on  water 
at  100°  only,  and  are  much  assisted  by  contact  with  another  metal. 

In  all  cases  in  which  cold  or  boiling  water  is  employed,  the  hydro- 
gen of  the  water  is  not  completely  displaced.  The  metal  forms  an 
hydroxide,  such  as  sodium  hydroxide  or  magnesium  hydroxide: 

Na  +  H2O  ->  H  +  NaOH, 
Mg  +  2H20  -»  2H  +  Mg(OH)2. 

With  sodium  the  resulting  solution  has  a  soapy  feeling  and  turns 
litmus  from  red  to  blue,  a  color  reaction  which  is  the  precise  opposite 
of  that  of  acids  (p.  51).  Substances  causing  these  two  effects  are 
called  alkalies,  and  the  soapy  feeling  and  the  action  on  litmus  are 
tests  for  alkalies.  Evaporation  reveals  the  sodium  hydroxide,  the 
alkali,  as  a  white  solid. 

With  steam  at  a  red  heat,  metals  like  iron,  zinc,  and  magnesium 
interact  vigorously.  The  steam,  generated  in  a  flask,  enters  at  one 
end  of  the  tube  containing  the  metal  (Fig.  23),  and  the  hydrogen 
passes  off  at  the  other.  Since,  at  a  red  heat,  all  hydroxides,  except 
those  of  potassium  and  sodium,  are  decomposed  into  an  oxide  of  the 


HYDROGEN 


6T 


metal  and  water,  as,  for  example,  Mg(OH)2  — » MgO  -f  H2O,  the  oxides 
are  formed  in  this  case: 

Mg  +  H2O  ->  MgO  +  2H. 
Iron  gives  the  magnetic  oxide,  Fe3O4.     We  note  that,  to  make  this 


FIG.  23. 


substance  (pp.  41-42),  four  unit-  weights  of  oxygen,  and  therefore 
four  formula-  weights  of  water,  are  required: 


4H2O 


3Fe  ->  Fe3O4  +  8H. 


The  Other  Ways  of  Preparing  Hydrogen.  —  For  special 
purposes,  hydrogen  may  be  made  by  boiling  an  aqueous  solution  of 
sodium  hydroxide  with  aluminium  turnings,  when  sodium  aluminate 
is  formed:  Al  +  NaOH  +  H2O  -*  NaA102  +  3H;  also  by  heating 
powdered  zinc  and  dry  sodium  hydroxide,  the  product  being  sodium 
zincate:  Zn  +  2NaOH  -»  Na2ZnO2  +  2H. 

Preparation  of  Simple  Substances.  —  There  are  two  general 
ways  of  obtaining  simple  substances,  both  of  which  have  now  been 
illustrated.  If  the  element  occurs  uncombined  in  nature,  as  oxygen, 
sulphur,  and  gold  do,  it  is  only  necessary  to  free  it  from  foreign 
materials  (impurities)  with  which  it  is  mixed.  If  no  such  supply 
exists,  as  in  the  case  of  hydrogen,  or  if  the  purification  is  difficult, 


68  COLLEGE   CHEMISTRY 

then  some  compound,  natural  or  artificial,  is  decomposed  and  the 
element  liberated. 

The  liberation,  in  turn,  may  be  effected  in  two  ways.  The  com- 
pound may  (1)  be  forced  apart  by  the  application  of  energy,  usually 
in  the  form  of  heat,  as  with  some  compounds  of  oxygen  (p.  46),  or 
electricity,  as  in  the  liberation  of  hydrogen  and  metals  by  elec- 
trolysis. Or  (2)  the  desired  constituent  may  be  liberated,  as  in 
the  case  of  hydrogen,  by  offering  to  the  other  constituents  some 
substance  with  which  they  will  unite. 

Displacement.  —  We  now  have  before  us  illustrations  of  two  sub- 
varieties  of  the  third  kind  (p.  10.  See  footnote  to  p.  24)  of  chemical 
change,  the  one  in  which  compounds  are  decomposed  and  the  parts 
combine  in  a  new  way.  The  first  sub-variety  was  double  decom- 
position, as  in  the  action  of  sodium  chloride  upon  silver  nitrate: 

NaCl  +  AgN03  -»  AgCl  +  NaNO3. 

In  that  class  of  cases  two  compounds  interact,  each  splits  into  the 
radicals  of  which  it  is  composed,  and  two  new  compounds  are  formed 
by  union  of  the  radicals  crosswise.  The  actions  just  used  in  the 
preparation  of  hydrogen  differ  from  these,  inasmuch  as  one  compound 
and  one  element  interact,  the  compound  splits  into  its  radicals,  and  one 
compound  and  one  free  element  are  produced: 

Zn  +  H2SO4  ->  ZnSO4  +  2H, 
Zn  +  2NaOH  ->  Na.>Zn02  +  2H. 

The  interacting  element,'  here  the  zinc,  is  said  to  displace  the  other, 
here  the  hydrogen,  from  combination.  In  double  decomposition 
there  is  an  even  exchange,  the  sodium,  for  example,  giving  up  one 
radical  (Cl),  and  getting  another  (NO3).  In  displacement  one  ele- 
ment gains  a  radical  while  another  loses  it,  the  zinc,  for  example, 
giving  up  nothing  but  getting  SO4,  while  the  hydrogen  loses  SO4,  and 
gains  nothing  in  return. 

Valence.*  —  We  shall  gain  much  help  in  the  making  of  equations 
if  we  now  introduce  and  bring  into  relation  to  the  symbols  a  concep- 
tion for  which  the  remarks  about  atomic  weights  (p.  36)  have  paved 
the  way.  It  will  have  been  observed  that  the  composition  of  the 

*  If  desired,  the  sections  on  Valence  may  be  taken  up  equally  well  after 
the  remaining  sections  of  this  chapter  have  been  studied. 


HYDROGEN  69 

chlorides  of  aluminium,  tin,  and  sodium  are  represented  by  the 
formulae  A1C13,  SnCl2,  and  NaCl,  respectively.  Again,  the  hydroxide 
of  sodium  is  NaOH,  while  those  of  magnesium  and  calcium  are 
Mg(OH)2  and  Ca(OH)2.  In  making  equations  we  constantly  need 
to  know  whether  the  chloride  of  an  element,  say  magnesium,  is 
MgCl,  or  MgCl2,  or  MgCl3,  or  MgCl4,  etc.,  and  whether  its  sulphate  is 
MgSO4,  or  Mg2S04,  or  some  other  combination  of  the  symbols.  To 
answer  questions  like  this  it  is  not  necessary  to  know  the  formula 
of  every  compound  of  each  element:  the  apparent  disorder  of  these 
numbers  can  be  reduced  to  rule,  and  the  reader  should  endeavor 
thoroughly  to  master  the  rule  before  going  farther. 

If  the  method  by  which  the  atomic  weights  were  derived  from 
equivalents  (p.  32)  is  now  reexamined,  the  nature  of  this  rule  will 
be  seen.  It  was  found,  for  example,  that  9.03  parts  of  aluminium 
(p.  36)  combined  with  the  equivalent  weights  of  the  other  elements, 
and  therefore  with  35.45  parts  of  chlorine.  If  this  weight  of  alumi- 
nium had  been  accepted  as  the  final  unit  (the  atomic  weight),  then 
it  would  have  been  represented  by  the  symbol  Al,  and,  since  Cl 
stands  for  35.45  parts  of  chlorine,  the  formula  of  the  chloride  would 
have  been  A1C1.  In  point  of  fact,  however,  a  number  three  times  as 
large  as  the  equivalent,  namely,  27.1,  was  chosen  as  the  atomic 
weight  of  aluminium,  and  the  symbol  Al  stands  for  this  triple  quan- 
tity. If  the  equivalent  of  chlorine  had  also  been  tripled  in  making 
its  atomic  weight,  the  amounts  represented  by  the  symbols  would  still 
have  been  chemically  equivalent,  and  the  formula  would  still  have 
been  A1C1.  But  the  equivalent  of  chlorine  was  left  unaltered. 
Hence,  to  get  the  equivalent  amounts  (i.e.,  the  actual  combining 
quantities)  of  the  two  elements,  we  must  have  3C1  with  1A1.  The 
formula  is  thus  A1C13.  Now,  it  is  evident  that  this  tripling  of  the 
equivalent  of  aluminium  will  affect  the  formula?  of  all  its  compounds. 
Whenever  it  is  combined  with  an  element  which,  like  chlorine,  has 
identical  equivalent  and  atomic  weights,  the  formula  of  the  com- 
pound will  be  of  the  form  A1X3.  In  accordance  with  this  we  have 
the  bromide  AlBr3.  In  making  the  formula?  of  compounds  of  alumi- 
nium, the  chief  thing  to  be  kept  in  mind,  therefore,  is  the  fact  that  its 
atomic  weight  contains  three  equivalents  and  always  combines  with 
three  equivalents  of  another  element.  This  fact  we  state  by  saying 
that  the  valence  of  the  atomic  weight  of  aluminium  is  three,  or  simply 
that  the  element  aluminium  is  trivalent. 


70  COLLEGE    CHEMISTRY 

Similarly,  the  equivalent  of  tin  is  59.5  and  its  atomic  weight  is 
119.  This  atomic  weight  therefore  contains  two  equivalents  of  tin 
and  combines  with  two  equivalents  of  any  other  element.  Hence, 
the  formula  of  a  compound  of  tin  with  an  element  of  the  chlorine 
class  will  be  SnX2.  Thus  tin  is  bivalent.  In  like  manner  the  equiva- 
lent of  sodium  is  23,  and  this  number  was  not  altered  in  making  the 
atomic  weight.  Hence,  the  symbol  Na  stands  for  one  equivalent, 
and  the  formula  of  the  compound  with  chlorine  is  NaCl.  Elements 
whose  atomic  weights  are  identical  with  their  equivalents  are 
described  as  univalent. 

Thus  the  valence  of  an  element  may  be  defined  as  the  number  of 
equivalent  weights  contained  in  its  atomic  weight.  Arithmetically 
it  is  the  integer  by  which  the  equivalent  weight  was  multiplied  in 
forming  the  atomic  weight.  The  above  explanation  shows  that  we 
may  define  the  valence  of  an  element  also  as  the  number  of  atomic 
weights  of  a  univalent  element,  with  which  its  atomic  weight  will 
combine. 

Sometimes  the  valence  is  indicated  in  the  symbol  thus:  Alin, 
Sn",  Na1,  Cl1,  Br1.  The  table  of  atomic  weights  (p.  36)  shows  the 
following  additional  cases:  O",  Cu",  S",  Hg",  H1,  Fen,  Mg",  CIV. 
With  the  help  of  this  list  the  formulae  of  compounds  may  easily  be 
made.  Thus,  oxygen  is  bivalent,  and  an  atomic  weight  of  oxygen, 
represented  by  O,  will  combine  with  two  atomic  weights  of  a  univa- 
lent element  as  in  OIIH2I  (water),  or  with  one  atomic  weight  of  a 
bivalent  element  as  in  O"Snn  (stannous  oxide),  O"Hgn  (mercuric 
oxide),  OnCu"  (cupric  oxide),  O"Mgn  (magnesium  oxide).  Again, 
carbon  being  quadrivalent,  the  atomic  weight  combines  with  four 
units  of  chlorine  and  of  hydrogen  in  CIVC\41  (carbon  tetrachloride) 
and  C^Hj1  (methane),  or  with  two  units  of  oxygen  and  of  sulphur 
in  CIV02"  (carbon  dioxide)  and  CIVS2"  (carbon  disulphide).  When 
it  combines  with  a  trivalent  element,  equal  numbers  of  equivalents 
of  each  element  must  be  used,  as  in  C3IVAl4m  (aluminium  carbide), 
where  C3  and  A14  contain  twelve  equivalents  each.  This  method, 
with  a  very  few  exceptions,  will  give  the  formula  of  all  compounds 
containing  only  two  elements  —  so-called  binary  compounds. 

The  Valence  of  Radicals.  —  The  valence  of  elements  can  easily 
be  determined  when  they  are  present  in  binary  combination.  This  is 
no  longer  the  case  when  more  than  two  elements  are  united  together. 


HYDROGEN  71 

A  study  of  chemical  changes  shows,  however,  that  even  here  the 
conception  of  valence  can  still  be  employed.  In  the  interaction  of 
zinc  with  dilute  sulphuric  acid: 

Zn  +  H2SO4  -»  ZnSO4  +  2H 

the  group  SO4  passes  as  a  whole  from  combination  with  2H  to  com- 
bination with  Zn.  Hence,  although  we  cannot  by  inspection  deter- 
mine the  valence  of  sulphur,  we  do  perceive  that  the  radical  SO4,  as 
a  whole,  must  be  bivalent.  It  occurs,  in  fact,  in  all  sulphates,  as 
Ag2S04,  MgS04,  and  A12(SO4)3,  and  in  the  interactions  of  these  sub- 
stances it  usually  passes  intact  from  one  state  of  combination  to 
another,  and  behaves  as  if  it  were  a  unit  of  a  single  element  of  valence 
two.  Again,  in  the  interaction  of  salt  with  silver  nitrate  (p.  40), 
we  observe  that  the  radical  N03  is  univalent.  Still  again,  the  com- 
positions of  the  compounds  CaCl2  and  Ca(OH)2  show  that  the  radical 
OH  (hydroxyl)  is  univalent.  The  formula  NaOH  leads  to  the  same 
conclusion. 

This  addition  to  our  ideas  enables  us  greatly  to  extend  the  list  of 
substances  of  which  we  can  write  the  formulae.  Thus,  the  hydroxides 
all  contain  (OH)1,  e.g.  A1III(OH)3I  (aluminium  hydroxide),  SnII(OH)21 
(stannous  hydroxide),  Cu^OHy  (cupric  hydroxide) .  The  nitrates 
all  contain  (NO,)1,  as:  H'(NO3)' (nitric  acid),  Mg"(NOa)ax  (magne- 
sium nitrate). 

It  is  to  preserve  the  identity  of  the  radicals  that  we  write  them 
in  brackets  and  place  the  coefficient  outside,  as  Ca(OH)2  and 
Mg(NO3)2,  instead  of  using  the  forms  CaO2H2,  MgN2O6,  and  so  forth. 
In  fact,  substances  which  commonly  interact  as  if  the  radicals  were 
single  elements,  we  regard  as  binary  compounds. 

In  writing  formulae  of  inorganic  compounds  we  usually  place  the 
positive  radical  (p.  64)  in  front  and  the  negative  radical  after  it. 

Sow  to  Ascertain  the  Valence  of  an  Element  or  Radical.  — 

The  above  shows  that  the  valence  of  one  element  or  radical  may 
always  be  ascertained  by  examination  of  the  formula  of  a  compound 
containing  another  element  or  radical  of  known  valence.  Thus, 
when  we  know  the  formula  of  sodium  iodide  to  be  Na1!,  or  that  of 
hydrogen  iodide  to  be  H'l,  we  infer  that  iodine  is  univalent.  The 
formula  of  silica  (sand)  Si02"  shows  silicon  to  be  quadrivalent,  and 
indicates  that  the  chloride  must  be  SiCl4.  Similarly  the  formula  of 


72  COLLEGE  CHEMISTRY 

calcium  carbonate  CanCO3  shows  that  the  radical  C03,  which  is 
common  to  all  carbonates,  must  be  bivalent.  Hence,  the  chemist 
does  not  memorize  the  valences  themselves;  he  recovers  the  valence 
of  an  element  or  radical,  when  needed,  by  recalling  the  formula  of  a  sub- 
stance containing  this  element  or  radical  in  combination  with  a  more 
familiar  element  or  radical,  such  as  Cl1  or  H1. 

It  is  absolutely  essential  that  correct  valences  should  be  used  in 
constructing  equations,  and,  at  first,  the  student  will  find  the  task 
by  no  means  easy.  He  should  give  special  attention  to  this  matter 
until,  by  solving  the  exercises  at  the  end  of  this  chapter,  and  by 
careful  examination  of  all  the  equations  encountered  in  the  text,  he 
has  mastered  the  subject. 

Multiple  Valence.  —  Some  elements  show  more  than  one  valence. 
This  is  as  much  as  to  say  that  an  atomic  weight  of  such  an  element 
may  form  stable  compounds  with  two,  or  even  more  different  num- 
bers of  equivalents  of  another  element.  This  fact  has  already  been 
mentioned,  for  it  is  implied  in  the  law  of  multiple  proportions  (p.  33). 
Thus  an  atomic  weight  of  tin  may  form  two  different  compounds  with 
chlorine,  namely,  Sn"Cl2  (st&nnous  chloride)  and  SnIVCl4  (stannic 
chloride).  Tin  behaves  in  the  same  way  towards  other  elements, 
however,  and  we  have  a  series  of  stannoi^s  compounds,  SnO,  SnBr2> 
and  so  forth,  and  a  corresponding  series  of  stanmc  compounds, 
SnO2,  SnBr4,  etc.  Two  different  valences  of  the  same  element  or 
radical  gives  rise  therefore  to  two  complete  sets  of  compounds.  The 
nomenclature  used  to  distinguish  the  two  series  has  been  discussed 
before  (p.  51).  As  a  rule,  an  element  passes  from  one  form  of  com- 
bination to  another  without  change  of  valence.  But  compounds  of 
elements  like  tin  can  also  undergo  changes  in  course  of  which  the 
valence  alters. 

Physical  Properties  of  Hydrogen.  —  Some  of  these  may  be 
given  in  tabular  form: 

Colorless  Boiling-point,  -  252.5° 

Tasteless  and  Odorless  Melting-point  (58  mm.),  -  260° 

Wt.  of  1  1.,  0.08987  g.  Sol'ty  in  Aq,  1.9  vote,  in  100  (14°) 

Air  is  14.5  times  as  heavy,  hence  the  gas  may  be  poured  upwards 
and  is  used  for  filling  balloons.  A  liter  flask  filled  with  air  requires 


HYDROGEN 


73 


about  1.2  g.  to  be  added  to  the  tare  to  restore  the  balance  when  the 
air  is  displaced  by  hydrogen.  Hydrogen  was  first  liquefied  in  visible 
amounts  by  Dewar  (1898).  The  liquid  is  colorless,  and,  when 
allowed  to  evaporate  rapidly  under  reduced  pressure,  freezes  to  a 
colorless  solid.  All  other  gases  except  helium  (q.v.)  solidify  easily 
when  led  into  a  vessel  surrounded  by  liquid  hydrogen. 

Hydrogen  is  absorbed,  for  the  most  part  in  a  purely  mechanical 
way,  by  many  metals.  Heated  iron  will  take  up  19  times  its  volume 
of  hydrogen,  gold  takes  up  46  volumes,  platinum  in  fine  powder  50 
volumes,  palladium  502  volumes,  and  silver  none.  The  maximum 
absorbed  by  palladium  under  favorable  conditions  is  873  volumes. 

Diffusion.  —  If  a  volume  of  gas  is  inclosed  at  one  end  of  a  cylin- 
der, the  rest  of  which  is  entirely  empty,  and  is  suddenly  released 
from  this  confinement,  it  spreads  with  extreme  speed 
so  as  to  occupy  the  whole  of  the  cylinder  to  an  equal 
degree.  This  spreading  is  not  an  effect  of  gravitation, 
since  it  takes  place  upwards  or  downwards  with  equal 
celerity.  The  same  phenomenon  is  observed  when, 
in  everyday  life,  a  bottle  of  scent  is  opened.  The 
vapor,  on  escaping,  begins  to  penetrate  in  all  direc- 
tions through  the  room,  showing  its  presence  by  its 
odor.  The  motion,  as  this  instance  shows,  takes 
place  through  a  space  occupied  by  another  gas  more 
slowly  than,  but  just  as  surely  as,  when  the  space 
is  empty.  The  material  of  gases  has  in  fact  an 
independent  power  of  locomotion.  The  resulting 
phenomenon  we  call  diffusion.  It  is  constant  in  rate 
for  each  gas  under  like  conditions,  and  hydrogen  has 
the  greatest  speed  of  diffusion  of  all  the  gases. 

The  different  rates  of  diffusion  of  different  gases 
are  easily  shown  by  comparing  their  several  speeds 
with  that  of  air,  when  both  pass  through  a  wall  of  unglazed,  porous 
porcelain. 

The  porous  cylinder  A  (Fig.  24)  contains  air  and  is  connected  with 
a  wide  tube  which  dips  beneath  the  surface  of  the  water.  When  a 
cylinder  H  containing  hydrogen  is  brought  over  it,  rapid  escape  of 
gas  takes  place  through  the  water,  showing  that  a  rise  in  pressure  has 
taken  place  inside  the  porous  vessel.  Before  the  cylinder  of  hydro- 


74  COLLEGE   CHEMISTRY 

gen  approached  it,  the  air  was  moving  both  outwards  and  inwards 
through  the  porcelain,  but,  being  the  same  air,  the  speed  of  motion 
was  equal  in  both  directions,  and  therefore  the  pressure  inside  was 
not  affected.  It  is  important  to  note  that  there  was  at  no  time  rest, 
there  was  simply  equal  motion  in  both  directions.  When  the  hydro- 
gen atmosphere  surrounded  the  cylinder,  the  hydrogen  gas  moved 
more  rapidly  into  the  cylinder  than  the  air  inside  could  move  out, 
and  hence  an  excess  of  pressure  quickly  arose  in  the  interior. 

Exact  measurement  shows  that  the  lighter  a  gas  is  in  bulk,  the 
faster  its  parts  move  by  diffusion  in  any  direction.  The  rate  is 
inversely  proportional  to  the  square  root  of  the  density  of  the  gas. 
Thus,  for  hydrogen  and  air  it  is  in  the  ratio  \/l  :  V.0695,  or  3.8  :  1. 

Chemical  Properties  of  Hydrogen.  —  Hydrogen,  delivered 
from  a  jet,  burns  in  air  or  pure  oxygen.  A  cold  vessel  held  over  the 
almost  invisible  blue  flame  condenses  to  droplets  of  water  the  steam 
that  is  produced.  Although  the  flame  gives  little  light,  it  is  exceed- 
ingly hot.  Platinum  melts  in  it  easily.  In  a  closed  space  it  pro- 
duces a  temperature  of  over  2500°.  When  hydrogen  and  oxygen 
are  mingled  in  a  suitable  burner,  and  the  flame  is  allowed  to  play 
on  a  piece  of  quicklime,  the  latter  becomes  white-hot  at  the  spot 
where  the  flame  meets  it.  This  result  is  called  a  calcium  light  or 
lime  light. 

When  hydrogen  and  oxygen  are  mixed,  the  chemical  action  is 
very  slow  at  ordinary  temperatures,  no  perceptible  amount  of  union 
occurring  in  a  period  of  five  years.  If  the  mixture  is  sealed  up  and 
kept  at  300°,  after  several  days  a  small  part  is  found  to  have  com- 
bined to  form  water.  At  518°,  hours  are  required  before  the  union 
is  complete.  At  700°  the  combination  is  almost  instantaneous. 
Hence  contact  with  a  body  at  a  bright-red  heat  (p.  53)  is  required 
actually  to  explode  the  mixture. 

Finely  divided  platinum,  when  held  in  the  cold  mixture,  hastens 
the  union  (otherwise  vanishingly  slow)  in  the  part  of  the  gases  in 
contact  with  it.  The  heat  of  the  union  raises  the  temperature  of 
neighboring  portions  and  causes  explosion  of  the  mass.  The  plati- 
num is  simply  a  catalytic  agent  (p.  54)  and  remains  itself  unaffected. 

Hydrogen  unites  directly  with  a  minority  only  of  the  simple  sub- 
stances. It  combines  rapidly  with  oxygen,  chlorine,  fluorine,  and 
lithium,  and  more  slowly  with  a  few  others. 


HYDROGEN  75 

When  these  elements,  especially  the  first  two,  are  already  in  com- 
bination, hydrogen  may  still  sometimes  displace  the  material  with 
which  they  are  united.  Thus,  when  any  one  of  the  oxides  of  iron  is 
heated  in  a  tube  through  which  hydrogen  flows,  the  latter  combines 
with  the  oxygen  to  form  water,  and  the  metal  is  liberated.  The 
skeleton  equation  (p.  42)  is:  Fe3O4  +  H  — >  H2O  +  Fe.  We  then 
reason  that  Fe3  will  give  3Fe.  Since  all  the  oxygen  is  removed 
from  the  compound,  O4  will  give  4H2O.  To  produce  this,  8H  is 
required.  Hence : 

Fe3O4  +  8H  ->  4H2O  +  3Fe. 

This  interaction  is  classed  as  a  displacement.  In  describing  it  the 
chemist  would  also  say  that  the  hydrogen  has  been  oxidized  and  that 
the  oxide  of  the  metal  has  been  reduced  (pp.  52-53). 

Specific  Chemical  Properties.  —  If  the  foregoing  section  on  the 
chemical  properties  of  hydrogen  and  the  corresponding  section  under 
oxygen  (p.  48)  are  now  reexamined,  the  nature  of  the  facts  contained 
in,  and  to  be  learned  from  such  a  section  will  be  seen.  Under  this 
head  we  describe  the  chemical  behavior  of  a  substance,  (1)  enumerat- 
ing the  other  substances,  simple  or  compound,  with  which  it  unites 
or  interacts,  (2)  stating  the  conditions  peculiar  to  each  action,  and 
(3)  estimating  the  intensity  of  the  tendency  to  chemical  change  in 
each  case.  In  the  case  of  a  simple  substance  like  oxygen  we  note 
particularly  with  how  many  of  the  other  elements  it  can  form  com- 
pounds, how  far  it  unites  with  them  directly,  and  in  how  many  cases 
the  compounds  have  to  be  made  by  indirect  means.  In  general, 
we  call  those  simple  substances  active  which  unite  with  many  other 
simple  substances  and  do  so  by  direct  union.  Oxygen,  for  example, 
is  active,  and  nitrogen  (q.v.)  is  relatively  inert. 

The  Speed  of  Chemical  Actions ;  a  Means  of  Measuring  Ac- 
tivity. —  One  means  of  measuring  the  relative  chemical  activi- 
ties of  several  substances  is  to  observe  the  speed  with  which  they 
undergo  the  same  chemical  change  (p.  19).  Thus  we  may  compare 
the  activities  of  the  various  metals  by  allowing  them  separately  to 
interact  with  hydrochloric  acid  and  collecting  and  measuring  the 
hydrogen  liberated  per  minute  by  each.  It  will  be  seen,  even  in  the 
roughest  experiment,  that  magnesium  is  thus  much  more  active  than 


76  COLLEGE   CHEMISTRY 

zinc.  The  comparison  must  be  made  with  such  precautions,  how- 
ever, as  will  make  it  certain  that  the  conditions  under  which  the 
several  metals  act  are  all  alike.  Thus,  in  spite  of  the  heat  evolved 
by  the  action,  means  must  be  used,  by  suitable  cooling,  to  keep  the 
temperature  at  some  fixed  point  during  the  experiment,  for  all 
actions  become  more  rapid  when  the  temperature  rises  (p.  53). 
Again,  the  pieces  of  the  various  metals  must  be  arranged  so  that 
equal  surfaces  are  exposed  to  the  acid  in  each  case.  It  is  found  that 
the  order  in  which  this  comparison  places  the  metals  is  much  the 
same  as  that  in  which  they  are  placed  by  a  study  of  other  similar 
actions.  A  single  table  suffices,  therefore,  for  all  purposes  (see 
Electromotive  series,  also  footnote  to  p.  24). 

Exercises.  —  1.  What  are  the  valences  of  the  negative  radicals 
of  phosphoric  acid  (p.  51),  and  of  acetic  acid  (p.  64)?  What  must 
be  the  formulae  of  calcium  phosphate,  cupric  acetate,  aluminium 
phosphate,  ferrous  carbonate,  ferrous  sulphate,  cupric  chloride? 

2.  What  is  the  valence  of  phosphorus  in  phosphoric  anhydride 
(p.  51)?     What  must  be  the  formulae  of,    (a)   the  corresponding 
chloride  and  sulphide  of  phosphorus,  and  (6)  of  aluminium  oxide? 

3.  What  are  the  valences  of  the  elements  in  the  following:  LiH, 
NH3,  SeH2,  BN? 

4.  What  are  the  valences  of  the  metals  and  radicals  in  the  follow- 
ing: Pb(NO3)2,  Ce(SO4)2,  KC1,  KMnO4  (potassium  permanganate)? 
Name  all  the  substances  in  3  and  4. 

5.  Write  the  formulae  of  ferrous  and  ferric  oxides,  of  ferrous  and 
ferric  nitrates,  of  stannous  and  stannic  sulphides. 

6.  Make  equations  to  represent,  (a)  the  reduction  of  lead  dioxide 
(Pb02)  by  hydrogen,  (b)  the  actions  of  aluminium  upon  cold  water 
and,  (c)  upon  steam  at  a  red  heat. 


CHAPTER  VIII 
WATER 

THE  great  quantity  of  water  which  occurs  in  nature  makes  it  one 
of  the  most  familiar  chemical  substances.  Water  is  found  also  in  the 
bodies  of  both  animals  and  plants  in  large  amounts,  and  is  indeed 
essential  to  the  working  of  living  organisms. 

Natural  Waters.  —  Sea-water  holds  about  3.6  per  cent  of  solid 
matter  in  solution.  Rain-water  is  the  purest  natural  water,  but 
contains  nitrogen,  oxygen,  and  carbon  dioxide  dissolved  from  the 
air.  Well-waters  often  contain  calcium  sulphate,  calcium  bicarbo- 
nate, and  compounds  of  magnesium  in  solution,  and  are  then  described 
as  hard.  Other  waters  contain  compounds  of  iron,  and  still  others 
are  effervescent  and  give  off  carbon  dioxide.  These  are  called  min- 
eral waters.  All  of  the  dissolved  substances  are  obtained  by  the 
water  in  its  progress  over  or  under  the  surface  of  the  ground. 

Water  which  is  to  be  used  for  domestic  purposes  is  examined,  par- 
ticularly, for  organic  matter.  This  usually  gains  access  to  the  water 
by  admixture  of  sewage  (p.  52).  It  is  not  this  organic  matter 
itself  which  is  deleterious,  but  the  bacteria  of  putrefaction  and  disease 
and  their  products  (ptomaines)  which  are  likely  to  accompany  it. 

Purification  of  Water. —  The  foreign  materials  which  water  may 
contain  are  divisible  into  two  kinds,  —  dissolved  matter  and  sus- 
pended matter.  No  water  is  free  from  either  of  these  varieties  of 
impurity.  In  chemical  laboratories  distilled  water  is  always  em- 
ployed. Yet  it  is  difficult  to  keep  the  liquid  pure  even  for  a  short 
time.  Ordinary  glass  dissolves  in  water  to  a  very  noticeable  extent. 

For  ordinary  purposes  the  suspended  matter  which  water  contains 
is  removed  by  nitration  (p.  7).  In  the  laboratory  this  takes  place 
through  unsized  paper.  The  pores  of  the  paper  are  sufficiently  small 
to  retain  suspended  particles,  while  permitting  the  passage  of  the 
water  with  its  dissolved  matter.  On  a  large  scale,  beds  of  gravel  are 

77 


78  COLLEGE   CHEMISTRY 

employed.  In  the  household,  the  Pasteur  filter  is  more  compact  and 
efficient.  The  water  is  forced  by  its  own  pressure  through  the  pores 
of  a  closed  tube  made  of  unglazed  porcelain.  Care  must  be  taken 
to  clean  these  tubes  at  frequent  intervals,  so  that  organic  and  per- 
haps putrescent  matters  may  not  accumulate  upon  them. 

Matter  in  solution  cannot  be  removed  by  ordinary  filtration,  and 
is  eliminated  by  distillation  (p.  26).  Since  the  water  is  converted 
into  steam  and  is  condensed  in  platinum  or  tin  pipes,  only  gases  or 
volatile  liquids  dissolved  in  it  can  pass  into  the  distillate. 

Physical  Properties  of  Water.  —  A  deep  layer  of  water  has  a 
blue  or  greenish-blue  color.  At  a  pressure  of  760  mm.,  water  ex- 
ists as  a  liquid  between  0°  and  100°.  Below  0°  it  becomes  solid, 
above  100°  a  gas.  Of  all  chemical  substances  it  is  the  one  which  we 
use  most,  so  that  its  physical  properties,  discussed  below,  should 
be  studied  attentively.  Then  too,  what  is  said  of  water  is  in  gen- 
eral true  of  all  other  liquids,  from  which  it  differs  only  in  details. 

Ice.  —  The  raising  or  lowering  of  the  temperature  of  a  gram  of 
water  through  one  degree  involves  the  addition  or  removal  of  one 
calorie  of  heat.  The  conversion,  however,  of  a  gram  of  water  at 
0°  to  a  gram  of  ice  at  0°  requires  the  removal  of  79  calories.  The 
mere  melting  of  a  gram  of  ice  causes  an  absorption  of  heat  to  the 
same  amount,  called  the  heat  of  fusion  of  ice.  At  0°  a  mixture  of 
ice  and  water  will  remain  in  unchanged  proportions  indefinitely. 
Any  cause  which  tends  permanently  to  lower  or  raise  the  temperature 
by  a  fraction  of  a  degree,  however,  will  bring  about  the  disappearance 
of  the  water  or  of  the  ice  respectively.  This  temperature  is  called 
the  melting  or  the  freezing  point.  A  temperature,  like  this,  at  which 
a  substance  passes  from  one  physical  state  of  aggregation  to  another  is 
called  a  transition  point. 

Steam  and  Aqueous  Tension.  —  At  atmospheric  pressure,  water 
passes  into  steam  rapidly  at  100°,  but  at  lower  temperatures,  and 
even  when  frozen,  it  does  the  same  thing  more  slowly.  The  quantity 
of  the  vapor  present  is  defined  by  the  gaseous  pressure  it  exercises, 
the  value  being  called  the  vapor  pressure  of  water  vapor  (or  of  the 
vapor  of  any  other  volatile  substance)  in  the  location  in  question. 

The  most  significant  fact  about  vapor  pressure  is  that,  when  excess 
of  the  liquid  is  present,  the  pressure  of  the  vapor  quickly  reaches  a 


WATER 


79 


definite  maximum  value  for  each  temperature.  In  the  absence  of 
excess  of  the  water,  less  than  this  maximum  pressure  may  exist. 
More  than  the  maximum  pressure  proper  to  a  given  temperature,  if 
produced  by  compression,  cannot  be  maintained,  for  a  part  of  the 
vapor  condenses  to  the  liquid  state.  The  magnitude  of  this  maxi- 
mum vapor  pressure,  at  a  given  temperature,  depends  on  the  ability 
of  the  particular  liquid  to  generate  vapor.  This  maximum  vapor 
pressure  is  held,  therefore,  to  represent  the  vapor  tension  of  the  liquid, 
at  the  given  temperature,  and  this  is  a  specific  property  of  the 
substance. 

I/  The  vapor  tension  may  be  shown  by  allowing  a  few  drops  of  water 
to  ascend  into  a  barometric  vacuum  (Fig.  25).  The  tube  on  the  left 
shows  the  mercury  when  nothing  presses  on  its 
surface.  The  tube  on  the  right  shows  the  result 
of  admitting  the  water.  The  difference  in  the 
height  of  the  two  columns  gives  the  value  of 
the  vapor  pressure  of  the  water  vapor.  With 
excess  of  water,  the  value  is  that  of  the  vapor 
tension,  called,  in  the  case  of  water,  the  aqueous 
tension.  The  jacket  surrounding  the  tube  on  the 
right  enables  us,  by  adding  ice  or  warm  water, 
to  maintain  any  temperature  between  0°  and 
100°.  When  ice  is  used  outside,  and  a  piece  of 
it  is  introduced  into  the  vacuum,  the  vapor  it 
gives  off  quickly  reaches  a  pressure  of  4.5  mm. 
The  vapor  pressure  of  the  ice  takes  the  place  of 
4.5  mm.  of  mercury  in  balancing  the  atmospheric 
pressure,  and  so  the  mercury  column  falls  by  this 
amount.  Similarly,  water  at  10°  causes  a  fall  of 
9.1  mm.  and  at  20°  of  17.4  mm.,  so  that  these 
represent  the  mercury-height  values  of  the 
aqueous  tension  at  these  temperatures.  The 
quantity  of  water  used  makes  no  difference,  so 
long  as  a  little  more  is  present  than  is  required 
to  fill  the  available  space  with  vapor.  With 
ether,  instead  of  water,  at  10°  the  fall  is  28.7  mm. 
With  water  at  higher  temperatures  the  fall  of 
the  mercury  column  becomes  much  greater.  At  50°  it  is  92  mm., 
at  70°  it  is  233.3  mm.,  at  90°  it  is  525.5  mm.,  and  at  100°  it  is  760 


FIG.  25. 


80  COLLEGE   CHEMISTRY 

mm.,  or  one  atmosphere.  At  121°  the  aqueous  tension  is  two 
atmospheres,  at  180°  it  is  ten  atmospheres. 

When  water  at  a  certain  temperature  has  given  the  full  amount  of 
water  vapor  to  the  space  above  it  that  its  aqueous  tension  permits, 
we  say  that  the  space  is  saturated  with  vapor.  That  concentration  of 
vapor  which  constitutes  saturation  varies  with  the  temperature  of 
the  water  and  depends  therefore  solely  on  the  power  of  the  water  to 
give  off  vapor.  It  has  nothing  to  do  with  the  size  of  the  space,  and 
is  even  independent  of  other  gases  the  space  may  already  contain 
(p.  60,  also  pp.  90-91.  See  footnote  to  p.  24). 

The  space  immediately  above  the  surface  of  the  ground,  which  is 
mainly  occupied  by  atmospheric  air,  is,  on  an  average,  less  than  two- 
thirds  saturated  with  water  vapor.  That  is  to  say,  such  air,  when 
inclosed  in  a  vessel  containing  water,  will  take  up  about  one-half 
more  than  it  already  contains.  The  vapor  of  water  at  100°  in 
an  open  vessel  displaces  the  air  entirely,  and,  if  the  required  heat  is 
furnished,  the  liquid  boils.  This  temperature,  like  the  freezing-point, 
is  a  transition  point. 

A  gram  of  water  at  100°,  in  turning  into  a  gram  of  steam  at  100°, 
takes  up  537  calories.  This  is  called  its  heat  of  vaporization.  Steam, 
in  fact,  contains  much  more  internal  energy  than  an  equal  weight 
of  water  at  the  same  temperature,  just  as  water,  in  turn,  contains 
more  energy  than  ice. 

All  our  substances  and  apparatus  have  traces  of  water,  derived 
from  the  atmosphere,  condensed  on  their  surfaces.  This  water  is, 
in  a  sense,  in  an  abnormal  condition,  for  it  does  not  evaporate  even 
in  dry  air.  It  is  observed  to  pass  off  in  vapor,  however,  when  we 
have  occasion  to  heat  the  substance  or  apparatus. 

Water  as  a  Solvent.  —  One  of  those  physical  properties  of  water 
which  are  most  used  in  chemical  work  is  its  tendency  to  dissolve 
many  substances.  This  subject  is  so  important  and  extensive  that 
we  shall  presently  devote  a  complete  chapter  to  some  of  its  simpler 
and  more  familiar  aspects. 

Chemical  Properties  of  Water.  —  Water  is  so  very  frequently 
used  in  chemical  experiments  in  which  it  is  a  mere  mechanical  ad- 
junct, that  the  beginner  has  difficulty  in  distinguishing  the  cases  in 
which  it  has  itself  taken  part  in  the  chemical  interaction.  The  four 


WATER  81 

kinds  of  chemical  activity  which  it  shows  should  therefore  receive 
careful  notice: 

1.  Water  is  a  relatively  stable  substance. 

2.  It  combines  with  many  oxides,  forming  bases  or  acids. 

3.  It   combines  with  many  substances,   chiefly  salts,   forming 
hydrates. 

4.  It  interacts  with  some  substances  in  a  way  described  as  hydro- 
lysis.   This  property  will  not  be  discussed  until  a  characteristic  case 
is  encountered. 

Water  a  Stable  Compound:  Dissociation.  In  the  case  of  a 
compound,  the  first  chemical  property  to  be  given  is  always,  whether 
the  substance  is  relatively  stable  or  unstable.  Usually  the  specifica- 
tion is  in  terms  of  the  temperature  required  to  produce  noticeable 
decomposition.  Thus,  potassium  chlorate  gives  off  oxygen  at  a  low 
red  heat.  Now,  water  vapor,  when  heated,  is  progressively  decom- 
posed into  hydrogen  and  oxygen,  yet  at  2000°  the  decomposition 
reaches  only  1.8  per  cent,  and  reunion  occurs  gradually  as  the  tem- 
perature is  lowered. 

A  decomposition  which  thus  proceeds  at  higher  temperatures, 
while  at  lower  temperatures  combination  of  the  constituents  can 
take  place,  is  called  a  dissociation.  The  decomposition  of  potassium 
chlorate  (p.  47)  is  not  a  dissociation  because  it  is  not  reversible; 
oxygen  will  not  under  any  circumstances  reunite  with  potassium 
chloride. 

Union  of  Water  with  Oxides.  —  When  sodium  combines  with 
oxygen  under  certain  conditions  we  obtain  sodium  oxide  (Na^O). 
The  product  unites  violently  with  water  to  form  sodium  hydroxide: 

Na2O  +  H2O  -+  2NaOlT. 

The  slaking  of  quicklime  is  a  more  familiar  action  of  the  same  kind : 
CaO  +  H20  -»  Ca(OH)2. 

No  other  products  are  formed.  The  clouds  of  steam  produced  in  the 
second  instance  are  due  to  evaporation  of  a  part  of  the  water  by  the 
heat  produced  in  the  formation  of  calcium  hydroxide.  The  aqueous 
solutions  of  these  two  products  have  a  soapy  feeling,  and  turn  red 
litmus  blue,  and  the  substances  therefore  belong  to  the  class  of 
alkalies  (pp.  66,  51)  or  bases.  Very  many  hydroxides  which  are  of 


82 


COLLEGE    CHEMISTRY 


the  same  nature,  for  example  ferric  hydroxide  Fe(OH)3  and  tin 
hydroxide  Sn(OH)2,  are  formed  so  slowly  by  direct  union  of  the 
oxide  and  water  that  they  are  always  prepared  in  other  ways. 

Some  oxides,  although  they  unite  with  water,  give  products  of  an 
entirely  different  character.  Phosphoric  anhydride  and  sulphur 
dioxide  are  of  this  class  and,  as  we  have  seen  (p.  51),  yield  acids. 

These  two  classes  of  final  products  are  so  different  that  we  make 
the  distinction  the  basis  for  classification  of  the  elements  present  in 
the  original  oxides.  The  elements,  like  sodium  and  iron,  whose 
oxides  give  bases,  are  called  metallic  elements ;  those,  like  phosphorus, 
whose  oxides  give  acids,  are  called  non-metallic  elements.  The  dis- 
tinguishing words  are  selected  because  the  division  corresponds,  in  a 
general  way  at  least,  with  the  separation  into  two  sets  to  which 
merely  physical  examination  of  the  elementary  substances  would  lead. 

Hydrates.  —  Many  substances  when  dissolved  in  water  and 
recovered  by  spontaneous  evaporation  of  the  solvent  enter  into 
combination  with  the  liquid.  The  products, 
which  are  solids,  are  called  hydrates.  That 
they  are  regular  chemical  compounds  is 
shown  by  the  following  three  facts:  (1) 
These  compounds  show  definite  chemical 
composition  expressible  by  formula  in 
terms  of  chemical  unit  weights  of  the 
constituents.  (2)  Often  much  heat  is 
given  out  in  their  formation.  Thus,  in  the 
case  of  washing  soda,  the  decahydrate  of 
sodium  carbonate  (Na2CO3,  10H20),  the  heat  of  the  union 
(p.  55)  is  8800  cal.  (3)  The  hydrates  have  physical  properties 
entirely  different  from  those  of  their  components.  Thus,  cupric 
sulphate,  often  called  anhydrous  cupric  sulphate  to  distinguish 
it  from  the  compound  with  water,  is  a  white  substance  crystal- 
lizing in  shining,  colorless,  needle-like  prisms.  The  pentahydrate 
(blue-stone  or  blue  vitriol),  which  crystallizes  from  the  aqueous  solu- 
tion, is  blue  in  color,  and  forms  larger  but  much  less  symmetrical 
(asymmetric  or  triclinic)  crystals  (Fig.  26) : 

CuSO4  +  5H2O  <=±  CuS04,  5H20. 

The  chemical  properties  show  hydrates  to  be  relatively  unstable. 
When  heated,  the  hydrates,  as  a  rule,  lose  none  of  the  constituents  of 


FIG.  26. 


WATER  83 

the  original  compound,  but  only  the  water,  in  the  form  of  vapor. 
When  melted,  or  when  dissolved  in  water,  the  hydrates  are  disso- 
ciated (p.  81)  into  water  and  the  original  substance.  The  aqueous 
solutions  made  from  the  anhydrous  substances  and  from  the  hydrates 
have  identical  physical  and  chemical  properties.  Hence  the  cheaper 
of  the  two  forms  is  generally  purchased,  and  many  of  the  chemicals 
used  in  the  laboratory  are  in  the  form  of  hydrates.  In  consequence 
of  the  ease  with  which  hydrates  give  up  water  we  write  their  for- 
mulae (e.g.  CuSO4,  5H2O)  so  that  the  water  and  original  substance 
are  separate.  A  formula  thus  modified,  so  as  to  show  some  favorite 
mode  of  behavior  of  the  substance,  is  called  a  reaction  (p.  11)  for- 
mula. The  formula  HloCuS09,  which  would  be  equally  correct,  is 
never  employed,  because  its  use  would  disguise  the  relation  of  the 
substance  to  cupric  sulphate. 

The  less  stable  hydrates  dissociate  very  readily.  Thus  the  deca- 
hydrate  of  sodium  sulphate,  Na2SO4,10H2O  (Glauber's  salt),  loses 
all  the  water  it  contains  (effloresces)  when  simply  kept  in  an  open 
vessel.  When  kept  in  a  dosed  bottle,  a  very  little  of  it  loses  water, 
and  then  the  decomposition  ceases.  The  cause  of  this  we  discover 
when  a  crystal  of  the  hydrate  is  placed  above  mercury,  like  the  ice 
or  water  in  Fig.  25  (p.  79).  It  shows  a  definite  aqueous  tension.  At 
9°  the  value  of  this  is  5.5  mm.  As  its  temperature  is  raised,  the 
tension  increases.  When  the  temperature  is  lowered,  on  the  other 
hand,  the  tension  diminishes,  the  mercury  rises,  and  a  part  of  the 
water  enters  into  combination  again.  Different  hydrates  show 
different  aqueous  tensions  at  the  same  temperature.  For  example,  at 
30°,  that  of  water  itself  is  31.5  mm.,  strontium  chloride  (SrCl2,  6H2O) 
11.5  mm.,  cupric  sulphate  (CuSO4,  5H20)  12.5  mm.,  barium  chloride 
(BaCl2,  2H2O)  4  mm. 

In  view  of  these  facts,  we  perceive  that  loss  of  water  by  efflores- 
cence is  like  evaporation.  Those  hydrates  which,  like  Glauber's 
salt  and  washing  soda,  have  a  vapor  tension  approaching  that  of 
water  itself,  lose  their  water  at  ordinary  temperatures  at  a  rapid 
pace.  Now,  atmospheric  air  is  usually  less  than  two-thirds  saturated 
with  water  vapor,  and  the  partial  pressure  (p.  60)  of  this  vapor 
opposes  the  dissociation  and  tends  to  prevent  the  liberation  of  the 
water.  Thus  at  9°,  the  vapor  tension  of  water  being  8.6  mm.,  the 
average  vapor  pressure  of  water  in  the  atmosphere  will  be  about 
5  mm.  Any  hydrate  with  a  greater  aqueous  tension  than  5  mm., 


84  COLLEGE   CHEMISTRY 

at  9°,  such  as  Glauber's  salt,  will  therefore  decompose  spontaneously 
in  an  open  vessel.  But  those  with  a  lower  vapor  tension,  such  as  the 
pentahydrate  of  cupric  sulphate  with  a  tension  of  2  mm.  at  9°,  will 
not  do  so  (see  pp.  90-93). 

The  water  of  hydration  is  known  colloquially  in  chemistry  as 
water  of  crystallization.  The  term  was  introduced  when  it  was  first 
observed  that  a  hydrate,  in  decomposing,  crumbles  and  loses  its 
original  crystalline  form.  But  the  term  is  misleading.  All  pure 
chemical  substances,  in  solid  form,  when  in  stable  physical  condi- 
tion, are  crystalline.  Amorphous  (i.e.,  non-crystalline)  substances, 
like  wax  and  glass,  are  supercooled  liquids. 

Composition  of  Water.  —  The  proportion  of  hydrogen  to  oxy- 
gen, in  water,  by  weight,  is  2  :  15.879  or  2.015  :  16.  The  proportion 
by  volume  is  2.0027  volumes  of  hydrogen  to  1  volume  of  oxygen. 
That  the  proportion  by  volume  is  very  close  to  2  :  1  may  easily 
be  shown  by  mixing  hydrogen  and  oxygen  in  this  proportion,  in  a 
strong  tube,  and  exploding  the  mixture  by  means  of  a  spark  from 
an  induction  coil.  The  resulting  steam  condenses  and  the  whole 
gas  vanishes.  If  different  proportions  are  used,  the  excess  of  one 
of  the  gases  remains  uncombined. 

Gay-Lussac's  Law  of  Combining  Volumes.  —  The  almost 
mathematical  exactness  with  which  small  integers  express  this  pro- 
portion is  not  a  mere  coincidence.  Whenever  gases  unite,  or  gaseous 
products  are  formed,  the  proportions  by  volume  (measured  at  the  same 
temperature  and  pressure)  of  all  the  gaseous  bodies  concerned  can  be 
represented  very  accurately  by  ratios  of  small  integers.  This  is  called 
Gay-Lussac's  law  of  combining  volumes  (1808).  Thus,  when  the  above 
experiment  is  carried  out  at  100°,  in  order  that  the  product,  water, 
may  be  gaseous  also,  it  is  found  that  the  three  volumes  of  the  con- 
stituents give  almost  exactly  two  volumes  of  steam.  For  example, 
15  c.c.  of  hydrogen  and  7.5  c.c.  of  oxygen  give  15  c.c.  of  steam.  Of 
course  the  hydrogen,  oxygen,  and  steam  must  be  measured  at  the 
same  pressure,  and  the  temperature  must  remain  constant  (100°) 
during  the  experiment.  Proper  manipulation  secures  the  former, 
and  a  jacket  filled  with  steam  (Fig.  27)  the  latter  condition.  Strips 
of  paper,  1,  2,  and  3,  are  pasted  on  the  jacket  in  such  a  way  that 
equal  lengths  of  the  eudiometer,  in  this  case  a  straight  one,  are  laid 
off.  The  three  divisions  being  filled  with  a  mixture  of  hydrogen 


WATER 


85 


and  oxygen  in  the  proper  proportions,  the  gas,  after  the  explosion, 
shrinks  so  as  to  occupy,  at  the  same  pressure,  only  two  of  them. 
From  this  universal  truth  in  regard  to  the  combination  of  gases, 
we  draw  the  important  inference  that  the  chemical  unit-weights  of 
simple  substances,  and  the  formula- weights 
of  compounds,  in  the  gaseous  condition, 
occupy  at  the  same  temperature  and  pressure 
volumes  which  are  equal,  or  else  stand  to 
one  another  in  the  ratio  of  small  integers. 

Exercises.  —  1.  Name  some  familiar 
transitions  (p.  78)  from  one  physical 
state  to  another. 

2.  What  evidence  is  there  in  the  com- 
mon behavior  of  ether  and  chloroform  to 
show  that  these  liquids  have  high  vapor 
tensions? 

3.  If  the  pressure  of  the  steam  in  a 
boiler  is  ten  atmospheres,  at  what  tem- 
perature is  the  water  boiling  (p.  80)  ? 

/  4.  How  many  grams  of  water  could  be 
heated  from  20°  to  100°  by  the  heat 
required  to  melt  1  kgm.  of  ice  at  0°? 

5.  What  do  you  infer  from  the  fact  that 
alum   and  washing  soda  lose  their  water 
of  hydration  when  left  in  open  vessels, 
while  gypsum  does  not? 

6.  Which  fact  shows  most  conclusively 
that  hydrates  are  true  chemical  compounds? 

7.  Gypsum  is  a  hydrate  of  calcium  sulphate  (CaS04).     If  6  g. 
of  gypsum,  when  heated,  lose  1.256  g.  of  water,  what  is  the  formula 
of  the  hydrate? 

8.  In  what  ways  does  a  hydrate  differ  from,  (a)  a  solution,  (6)  an 
hydroxide? 

9.  Should  you  expect  to  find  any  difference,  in  respect  to  chemical 
activity,  between  the  three  forms  of  water?     Have  we  had  any 
experimental  confirmation,  or  the  reverse,  of  this  conclusion  (p.  66)  ? 

10.  Name  some  crystalline  substances  which  are  not  used,  or  do 
not  occur  in  the  form  of  hydrates. 


FIG.  27. 


CHAPTER  IX 
THE    KINETIC-MOLECULAR    HYPOTHESIS 

As  soon  as  we  have  constructed  a  law  (p.  4)  we  desire  immediately 
to  find  out  the  basis  of  the  constant  mode  of  behavior  it  epitomizes. 
If  no  explanation,  that  is,  more  detailed  description,  is  forthcoming 
as  the  result  of  closer  observation,  we  proceed  to  imagine  one  (p.  5). 
This  always  takes  a  mechanical  form,  often  crude  at  first,  and  later 
undergoing  refinement.  Thus,  at  first,  the  phenomena  of  light  were 
explained  by  the  conception  of  clouds  of  fine  corpuscles  emanating 
from  the  luminous  body.  The  chances  of  hitting  upon  an  objective 
reality  by  guess-work  like  this  is  obviously  remote.  Whether  such 
particles  did  really  fly  about  was  not  the  main  question,  however. 
Their  value  lay  in  the  fact  that  they  could  be  pictured  concretely  and 
gave  a  basis  for  further  thought  and  perhaps  suggestions  for  new 
experiments.  Such  a  structure  of  the  imagination  is  called  an 
hypothesis. 

Tlie  Molecular  Hypothesis.  —  The  only  mechanical  basis  we  can 
imagine  to  account  for  the  physical  properties  of  matter  is  a  discon- 
tinuous structure  of  some  description.  The  fact  that  all  kinds  of 
matter  can  be  compressed  (gases  to  an  enormous  extent,  solids  and 
liquids  to  a  measurable  extent)  may  be  explained,  either  by  a  diminu- 
tion in  the  volume  of  the  material  itself,  or  by  the  closer  packing 
together  of  the  particles  into  which  this  material  is  divided.  It  is 
evident  that  the  latter  is  much  more  in  harmony  with  our  experience. 
Oxygen  at  760  mm.,  for  example,  can  be  reduced  by  pressure  to  one 
two-hundredth  of  its  volume,  or  even  less.  Compression,  we  imagine, 
therefore,  does  not  here  diminish  the  actual  volume  occupied  by 
oxygen,  but  crowds  the  particles  closer  together  and  diminishes  to 
one  two-hundredth  the  space  between  them.  Compressing  a  gas  is, 
in  fact,  mainly  compressing  the  empti/  space  of  which  we  imagine  it 
chiefly  to  consist.  The  same  hypothesis  will  furnish  a  concrete 
description  of  how  one  kind  of  matter  will  frequently  absorb  a  large 

86 


KINETIC-MOLECULAR   HYPOTHESIS  87 

quantity  of  another.  Thus  we  picture  the  hydrogen  gas  taken  up 
by  iron  and  other  metals  as  being  packed  away  in  the  spaces  between 
the  particles  of  the  metal.  So  also,  in  solution,  the  volume  of  the 
liquid  does  not  usually  increase  by  an  amount  equal  to  that  of  the 
substances  dissolved.  Hence  we  imagine  the  particles  as  possibly 
incompressible  and  the  interstices  between  them  as  furnishing,  a  part 
of  the  accommodation  for  the  foreign  material.  A  particle  is  a 
fragment  of  matter  which  can  be  seen  and  measured  directly,  and 
handled  separately.  Since  the  particles  of  this  hypothesis  have  none 
of  these  qualities,  we  distinguish  them  by  the  name  molecules.  Mole- 
cules are  the  imaginary  units  of  which  bodies  are  aggregates. 

In  the  following  sections  the  term  particle,  when  used  at  all,  is 
usually  equivalent  to  molecule,  and  is  employed  only  to  avoid  too 
frequent  occurrence  of  the  latter  word. 

Kinetic- Molecular  Hypothesis  Applied  to  Gases.  —  Let    us 

first  build  up  our  hypothesis  to  fit  the  qualitative  properties  of 
gases.  We  print  in  italics  each  fact,  and  in  black  type  the  part  of 
the  hypothesis  invented  to  fit  it.  The  most  remarkable  thing  about 
a  gas,  considering  the  looseness  with  which  its  material  is  packed,  is 
the  total  absence  in  it  of  any  tendency  to  settling  or  subsidence. 
Since  the  molecules  cannot  be  at  rest  upon  one  another,  as  the  great 
compressibility  shows,  we  are  driven  to  suppose  that  they  are  widely 
separated  from  one  another,  and  that  they  occupy  the  space,  otherwise 
a  complete  vacuum,  by  constantly  moving  about  in  all  directions.  But 
a  moving  aggregate  of  particles  which  does  not  even  finally  settle 
must  be  in  perpetual  motion.  We  must,  therefore,  imagine  the 
molecules  to  be  wholly  unlike  particles  of  matter  in  having  perfect 
elasticity,  in  consequence  of  which  they  undergo  no  loss  of  energy 
after  a  collision.  They  must  continually  strike  the  walls  of  the 
vessel  and  one  another  and  rebound,  yet  without  loss  of  motion. 
The  diffusibilitij  of  gases  and  their  mutual  permeability  require  no 
additional  assumptions.  The  fact  that  each  gas  is  homogeneous, 
efforts  to  sift  out  lighter  or  heavier  samples  having  failed,  requires 
the  supposition  that  all  the  molecules  of  a  pure  gas  are  closely  alike. 

Passing  now  to  Boyle's  law  (p.  59),  the  thing  to  be  accounted  for 
is  that  when  a  sample  of  a  gas  diminishes  in  volume,  its  pressure 
increases  in  the  same  proportion.  Let  the  diagram  (Fig.  28)  repre- 
sent a  cylinder  with  a  movable  piston,  upon  which  weights  may  be 


88  COLLEGE   CHEMISTRY 

placed  to  resist  the  pressure.  Now  the  pressure  exercised  by  the  gas 
under  the  piston  cannot  be  like  the  pressure  of  the  hand  upon  a  table, 
since  we  have  just  assumed  that  the  particles  are  not  even  approxi- 
mately at  rest,  and  the  spaces  between  them  are  enormous  compared 
with  the  size  of  the  molecules  themselves.  The  gaseous  pressure 
must  therefore  be  attributed  to  the  colossal  hailstorm  which  their 
innumerable  impacts  upon  the  piston  produce.  If  this  is  the  case, 
the  compressing  of  a  gas  must  consist  simply  in  moving  the  partition 
downwards  so  that  the  particles  as  they  fly  about  are 
gradually  restricted  to  a  smaller  and  smaller  space. 
Their  paths  become  on  an  average  shorter  and  shorter. 
Their  impacts  upon  the  wall  become  more  and  more 
frequent.  So  the  pressure  which  this  causes  becomes 

T         greater  and  greater,  and  is  proportional  to  the  degree 
of  crowding  (concentration)  of  the  molecules.    There  are 

two  other  points  to  be  added.  When  we  diminish  the 
volume  to  one-half,  we  find  from  experience  that  the 
pressure  becomes  exactly,  or  almost  exactly,  twice  as 


FIG.  28.  great.  This  must  mean  that  although  the  particles 
are  becoming  crowded  they  do  not  interfere  with  one 
another's  motion,  excepting  of  course  where  actual  collision  causes 
a  rebound.  Only  in  the  absence  of  interference  would  doubling 
the  number  of  molecules  per  unit  of  volume  give  exactly  double  the 
number  of  impacts  on  the  walls.  Hence  the  molecules  must  have 
practically  no  tendency  to  cohesion.  Finally,  the  molecules  must 
be  supposed  to  move  in  straight  lines  between  collisions. 

Boyle's  law  therefore  adds  four  more  conceptions  to  our  molec- 
ular hypothesis,  namely,  that  the  impacts  of  the  particles  produce 
the  pressure,  that  the  crowding  of  the  molecules  represents  the  con- 
centration of  the  material  and  that  the  particles  move  in  straight  lines 
and  show  almost  no  cohesion,  since  pressure  and  concentration  are  very 
closely  proportional  to  one  another. 

How,  now,  can  we  account  for  Charles'  law  (p.  60),  according  to 
which  an  increase  in  pressure  (or  in  volume)  results  from  heating  a 
mass  of  rapidly  moving  molecules?  The  action  of  a  particle  collid- 
ing with  a  surface  is  measured  in  physics  in  terms  of  its  mass  and  its 
velocity.  It  is  evident  that  heating  a  cloud  of  molecules  would  not 
increase  the  mass  of  each,  and  it  must  therefore  increase  the  velocity 
of  each  since  the  kinetic  energy  of  all  becomes  greater. 


KINETIC-MOLECULAR   HYPOTHESIS  89 

The  fact  that  the  combining  volumes  of  gaseous  substances  are  equal, 
or  stand  to  one  another  in  the  ratio  of  small  whole  numbers  (Gay- 
Lussac's  law,  pp.  84-85),  suggests  two  ideas:  First,  that  chemical 
combination,  considered  in  detail,  and  arranged  to  harmonize  with 
this  hypothesis,  would  involve  unions  of  a  few  particles  of  more  than 
one  kind  to  form  composite  molecules.*  And,  second,  that  a  simple 
integral  relation  must  be  assumed  to  exist  between  the  numbers  of 
molecules  in  equal  volumes  of  different  gases,  at  the  same  temperature 
and  pressure.  Avogadro  (1811),  the  professor  of  physics  in  Turin, 
put  forward  the  hypothesis  that  these  numbers  might  be  equal.  A 
more  strict  study  of  the  assumptions  we  have  been  making,  and  of 
some  additional  facts,  has  since  shown  that  no  other  conjecture  than 
Avogadro's  would  be  consistent  with  them.  Thus  it  now  bears  the 
relation  of  a  logical  deduction  from  the  kinetic-molecular  hypothesis 
and  the  properties  of  gases,  and  is  known  as  Avogadro's  hypothesis. 
It  may  also  be  put  in  the  form:  At  the  same  temperature  and 
pressure,  the  molecular  concentration  of  all  kinds  of  gases  has  the  same 
value. 

The  law  of  diffusion  (p.  74)  harmonizes  with  the  kinetic-molecular 
hypothesis  without  further  modification  of  the  latter. 

Finally,  gases  can  be  liquefied  by  sufficient  cooling  and  compression. 
This  fact  compels  us  to  suppose  that,  after  all,  even  gaseous  mole- 
cules have  a  tendency  to  cohesion.  This  cohesion  is  scarcely  percep- 
tible so  long  as  the  gas  is  warm  and  diffuse.  But  when  the  kinetic 
energy  of  the  molecules  is  sufficiently  reduced  by  cooling  (namely,  to 
the  critical  temperature),  and  the  molecules  are  brought  sufficiently 
close  together,  the  mutual  cohesion  of  the  molecules  causes  the  gas 
to  condense  and  assume  the  liquid  form.  The  critical  temperature 
of  oxygen  is  —  118°,  of  hydrogen  —  234°,  of  carbon  dioxide  31.35°, 
of  water  358°. 

We  may  summarize  the  facts  about  gases,  appearing  in  italics 
above,  with  the  corresponding  fictions,  in  heavy  type,  which  we 
have  added  one  by  one  in  manufacturing  our  hypothesis,  as 
follows : 

*  This  is  essentially  the  idea  used  by  Dalton,  before  Gay-Lussac's  law  was 
known,  however,  for  the  explanation  of  the  laws  of  chemical  combination. 
He  called  it  the  atomic  hypothesis  (q.v.). 


90 


COLLEGE    CHEMISTRY 


FACTS. 


HYPOTHESIS. 


Non-settling 
Comp  ressibility 
Diffusibility 
Permeability 

Homogeneity. 


Relation  of  pressure  and  concen- 
tration (Boyle's  law). 


Relation  of  volume  (or  pressure) 
and  temperature  (Charles'  law). 


Relation  of  atomic  weights  and 
volumes  (Gay-Lussac's  law). 


Law  of  diffusion. 
Gases  can  be  liquefied. 


'The  fictitious  particles  called  molecules 
are,  at  0°  and  760  mm.,  at  great  aver- 
age distances  from  one  another;  they 
are  in  constant  motion  and  have  per- 
fect elasticity. 

The  molecules  of  the  same  substance  are 
closely  alike. 

The  effect  of  pressure  is  produced  by  the 
impacts  of  the  molecules,  and  is  propor- 
tional to  the  degree  to  which  they  are 
crowded  together ;  the  molecules  move  in 
straight  lines,  and,  when  widely  scat- 
tered, exhibit  almost  no  tendency  to 
cohesion. 

A  rise  in  temperature  increases  the  velo- 
city and  therefore  the  kinetic  energy  of 
the  molecules. 

Chemical  union  consists  in  fusion  of  differ- 
ent kinds  of  molecules  (Dalton's  hypo- 
thesis), of  which  there  are  equal  num- 
bers in  equal  volumes  of  different  gases 
at  the  same  temperature  and  pressure 
(Avogadro's  hypothesis). 


Molecules  do  possess  a  tendency  to 
cohesion,  which  they  exhibit  when 
cooled  and  closely  crowded  together. 


Kinetic  Hypothesis  Applied  to  Liquids.  —  The  phenomena 
connected  with  surface  tension,  such  as  coherence  into  drops,  show 
that  cohesion  plays  a  larger  part  in  liquids  than  in  gases.  On  the 
other  hand,  liquids  which  are  capable  of  mixing  (e.g.  alcohol  and 
water),  when  placed  above  one  another  in  the  same  vessel,  do  mix, 
slowly,  by  diffusion.  This  indicates  that  motion  of  the  molecules, 
although  much  impeded  by  friction,  has  not  been  annihilated  by 
cohesion. 

The  formation  of  vapor  from  cold  liquids  likewise  requires  us  to 
make  the  same  supposition.  To  be  consistent,  we  have  also  to 
imagine  that  the  vapor  above  the  liquid,  for  example  the  water  in 
the  barometer  tube  in  Fig.  25  (p.  79),  is  not  composed  of  the  same 


KINETIC-MOLECULAR  HYPOTHESIS  91 

set  of  molecules  one  minute  as  it  was  during  the  preceding  minute. 
Their  motions  must  cause  many  of  them  to  plunge  into  the  liquid, 
while  others  emerge  and  take  their  places.  When  the  water  is  first 
introduced,  there  are  no  molecules  of  vapor  in  the  space  at  all,  so 
that  emission  from  the  water  predominates.  The  pressure  of  the 
vapor  increases  as  the  concentration  of  the  molecules  of  vapor 
becomes  greater,  hence  the  mercury  column  falls  steadily.  At  the 
same  time  the  number  of  gaseous  molecules  plunging  into  the 
water  per  second  must  increase  in  proportion  to  the  degree  to 
which  they  are  crowded  in  the  vapor.  The  rate  at  which  mole- 
cules return  to  the  water  thus  begins  at  zero,  and  increases  steadily : 
the  rate  at  which  molecules  leave  the  water  maintains  a  constant 
value.  Hence  the  rate  at  which  vapor  molecules  enter  the  water 
must  eventually  equal  that  at  which  other  molecules  leave  the  liquid. 
At  this  point,  occasion  for  visible  change  ceases  and  the  mercury 
comes  to  rest.  We  are  bound  to  think,  however,  of  the  exchange  as 
still  going  on,  since  nothing  has  occurred  to  stop  it.  The  condition 
is  not  one  of  rest  but  of  rapid  and  equal  exchange.  Such,  described 
in  terms  of  the  hypothesis,  is  the  state  of  affairs  which  is  character- 
istic of  a  condition  of  equilibrium.  The  condition  is  kinetic,  and 
not  static. 

Equilibrium.  —  This  term  is  used  so  often  in  chemistry,  and  is 
used  in  so  unfamiliar  a  sense,  that  the  reader  should  consider 
attentively  what  it  implies.  Three  things  are  characteristic  of  a 
state  of  equilibrium: 

1.  There  are  always  two  opposing  tendencies  which,  when  equi- 
librium is  reached,  balance  one  another.     In  the  foregoing  instance, 
one  of  these  is  the  hail  of  molecules  leaving  the  liquid,  which  is 
constant  throughout  the  experiment.     It  represents  the  vapor  ten- 
sion of  the  liquid  (p.  79).      The  other  is  the  hail  of  returning  molecules, 
which,  at  first,  increases  steadily  as  the  concentration  of  the  vapor 
becomes  greater.     This  is  the  vapor  pressure  of  the  vapor.     These 
have  the  effect  of  opposing  pressures,  and,  when  the  latter  becomes 
equal  to  the  former,  equilibrium  is  established.     In  all  cases  of  equi- 
librium we  shall  symbolize  the  two  opposing  tendencies  by  arrows, 
thus: 

Water  (liq.)  <=±  Water  (vapor). 

2.  Although  their  effects  thus  neutralize  one  another  at  equilib- 
rium, both  tendencies  are  still  in  full  operation.     In  the  case  in  point, 


92  COLLEGE   CHEMISTRY 

the  opposing  hails  of  molecules  are  still  at  work,  but  neither  can 
effect  any  visible  change  in  the  system.  Equilibrium  is  a  state  of 
balance  or  poise,  not  of  rest. 

3  (and  this  is  the  chief  mark  of  equilibrium).  A  slight  change  in 
the  conditions  produces,  never  a  great  or  sharp  change,  but  always, 
and  instantly,  a  corresponding  small  change  in  the  state  of  the  system. 
The  change  in  the  conditions  accomplishes  this  by  favoring  or  dis- 
favoring one  of  the  two  opposing  tendencies.  Thus,  for  example,  when 
the  temperature  of  a  liquid  is  raised,  the  kinetic  energy  of  its  mole- 
cules is  increased,  the  rate  at  which  they  leave  its  surface  becomes 
greater,  the  vapor  tension  increases  and,  hence,  a  greater  concen- 
tration of  vapor  can  be  maintained.  The  system,  therefore,  quickly 
reaches  a  new  state  of  equilibrium  in  which  a  higher  vapor  pressure 
exists  (p.  79).  A  heap  of  matter  on  a  table  is  not  in  equilibrium, 
because  addition  of  more  material  produces  no  response  until,  when 
a  very  great  quantity  is  added,  the  table  breaks.  But  a  body  on 
the  scales  is  in  equilibrium,  for  the  addition  of  the  smallest  particle 
produces  a  corresponding  inclination  of  the  beam. 

In  the  preceding  illustration,  the  evaporating  tendency  was 
favored  by  a  rise  in  temperature.  As  an  example  of  a  change  in 
conditions  disfavoring  one  tendency,  take  the  case  where  the  liquid 
is  placed  in  an  open,  shallow  vessel.  Here  the  condensing  tendency 
is  markedly  discouraged,  for  there  is  practically  no  return  of  the 
emitted  molecules.  Hence  complete  evaporation  takes  place.  Ele- 
vation of  the  temperature  hastens  the  process.  A  draft  insures  the 
total  prevention  of  all  returns,  and  has  therefore  the  same  effect.  The 
two  methods  of  assisting  the  displacement  of  an  equilibrium,  and 
particularly  the  second,  in  which  the  opposed  process  is  weakened 
and  the  forward  process  triumphs  solely  on  this  account,  should  be 
noted  carefully.  They  are  applied  with  surprising  effectiveness  in 
the  explanation  of  chemical  phenomena  (see  Chaps,  xi  and  xv). 

Kinetic  Hypothesis  Applied  to  Solids.  —  The  properties  of 
solids  differ  from  those  of  liquids  chiefly  in  the  fact  that  the  solid  has 
a  definite  form  of  which  it  can  be  deprived  only  with  difficulty. 
This  we  may  explain  in  accordance  with  the  kinetic  hypothesis  by  the 
supposition  that  the  cohesion  in  solids  is  very  much  more  prominent 
than  in  liquids.  We  obtain  solids  from  liquids  by  cooling  them;  in 
other  words,  by  diminishing  the  kinetic  energy  and  therefore  the 


KINETIC-MOLECULAR  HYPOTHESIS  93 

velocity  of  the  particles.  The  cohesive  tendency  of  the  latter  is  thus 
able  to  make  itself  felt  to  a  greater  extent.  If,  conversely,  we  heat  a 
solid,  or,  according  to  the  hypothesis,  if  we  increase  the  speed  with 
which  the  particles  move,  the  body  first  melts  and  gives  a  liquid, 
and  this  finally  boils  and  becomes  a  gas.  The  intrinsic  cohesion  of 
the  particular  substance  can  undergo  no  change,  but  the  increasing 
kinetic  energy  of  the  particles  steadily  and  continuously  obliterates 
its  effects.  Yet  some  motion  still  survives  in  a  solid.  Thus  we  find 
that  when  the  layer  of  silver  is  stripped  from  a  very  old  piece  of 
electroplate,  the  presence  of  this  metal  in  the  German  silver  or  copper 
basis  of  the  article  is  easily  demonstrated. 

The  tendency  of  all  solids  to  assume  crystalline  forms,  which  show 
definite  cleavage  and  other  evidences  of  structure,  distinguishes  them 
sharply  from  liquids.  The  force  of  cohesion  in  liquids  is  exercised 
equally  in  different  directions.  In  solids  it  must  differ  in  different 
directions  in  order  that  structure  may  result.  Since  each  substance 
shows  an  individual  structure  of  its  own,  these  directive  forces  must 
have  special  values  in  magnitude  and  direction  in  each  substance. 

A  crystal  arises  by  growth.  When  the  process  is  watched,  as  it 
occurs  in  a  melted  solid  or  an  evaporating  solution,  the  slow  and 
systematic  addition  of  the  material  in  lines  and  layers,  as  if  according 
to  a  regular  design,  is  one  of  the  most  beautiful  and  interesting  of 
natural  phenomena.  The  fern-like  patterns  produced  by  ice  on  a 
window-pane  show  the  general  appearance  characteristic  of  crystalli- 
zation in  a  thin  layer.  A  larger  mass  in  a  deep  vessel  gives  forms 
which  are  geometrically  more  perfect.  From  its  very  incipiency  the 
crystal  has  the  same  form  as  when,  later,  its  outlines  can  be  dis- 
tinguished by  the  eye.  Hence  the  outward  form  is  only  an  expres- 
sion of  a  specific  internal  structure  which  the  continual  reproduction 
of  the  same  outward  form  on  a  larger  and  larger  scale  leaves  as  a 
memorial  of  itself  in  the  interior. 

Crystal  Forms.  —  Crystalline  form  is  continually  used  in  iden- 
tifying (pp.  6,  9,  24-25)  the  substances  produced  in  chemical  actions. 
The  classification  of  crystalline  forms  is  carried  out  according  to 
the  degree  of  symmetry  of  the  crystals: 

1.  Regular  system.  5.   Monosymmetric,    or 

2.  Square  prismatic  system.  monoclinic  system. 

3.  Hexagonal  system.  6.   Asymmetric,    or 

4.  Rhombic  system.  triclinic  system. 


94 


COLLEGE    CHEMISTRY 


The  regular  system  presents  the  most  symmetrical  figures  of  all. 
Some  forms  which  commonly  occur  are  the  octahedron  (Fig.  29) 
shown  by  alum,  the  cube  (Fig.  6,  p.  8)  affected  by  common  salt, 
and  the  dodecahedron  (Fig.  30)  frequently  assumed  by  the  garnet. 


FIG.  29. 


FIG.  30. 


FIG.  31. 


The  square  prismatic  system  includes  less  symmetrical  forms  than 
the  previous  one,  since  the  crystals  are  lengthened  in  one  direction. 
Fig.  31  shows  the  condition  in  which  zircon  (ZrSiO4),  which  furnishes 
us  with  the  basis  of  certain  incandescent  illuminating  arrangements, 

occurs  in  nature.  The 
form  of  ordinary  hydrated 
nickel  sulphate  (NiS04, 
6H20)  is  similar  to  this. 

The  hexagonal  system, 
like  the  preceding,  fre- 
quently exhibits  elon- 
gated prismatic  forms, 


FIG.  32. 


FIG.  33. 


but  the  section  of  the  crystals  is  a  hexagon,  instead  of  a  square, 
and  the  termination  is  a  six-sided  pyramid.  Quartz  (Fig.  32), 
or  rock  crystal,  is  the  most  familiar  mineral  in  this  system.  Calcite 
(CaCO3),  which  is  chemically  identical  with  chalk,  or  marble,  takes 
forms  known  as  the  scalenohedron  (Fig.  33)  and  rhombohedron 
(Fig.  9,  p.  9),  which  are  classified  in  a  subdivision  of  this  system. 
Indeed,  recently  it  has  become  common  to  erect  this  into  a 


KINETIC-MOLECULAR   HYPOTHESIS 


95 


separate  system  (the  trigonal),  in  which  both  quartz  and  calcite 
are  included. 

The  rhombic  system  includes  the  natural  forms  of  the  topaz, 
and  of  sulphur  (Fig.  1,  p.  6),  as  well  as  that  of 
potassium  permanganate  (Fig.  34),  potassium 
nitrate  (Fig.  63),  and  many  other  substances. 
These  crystals  exhibit  a  good  deal  of  symmetry, 
but  their  section  is  always  rhombic,  and  hence 
the  name. 

The  monosymmetric  system  exhibits  forms  which 
have  but  one  plane  of  symmetry.  Gypsum  (Fig. 
35),  which  is  hydrated  calcium  sulphate  (CaSO4, 

2H,O),     and    felspar     are 

minerals   possessing  forms 

of     this     kind.       Tartaric 

acid,  rock  candy  (Fig.  36), 

potassium  chlorate,  and  hydrated   sodium 

carbonate    (washing  soda)   belong   to   this 

system. 

The    asymmetric    system    includes    forms 

which  have  no  plane  of  symmetry  whatever.  Blue  vitriol  (Fig. 
26,  p.  82),  CuS04,  5H20,  is  one  of  the  most  familiar  substances  of 
this  kind. 


FIG.  35. 


FIG.  36. 


CHAPTER  X 
SOLUTION 

SOLUTIONS  are  so  constantly  used  in  chemistry  that  some 
knowledge  of  their  properties  is  desirable  in  order  that  we  may 
employ  them  intelligently.  In  what  follows,  we  give  a  preliminary 
account  of  some  of  the  simpler  facts  about  solution. 

General  Properties  of  Solutions.  — •  A  solid  may  be  distributed 
through  a  liquid,  either  by  being  simply  suspended  (p.  77)  in  the 
latter  (mixture),  or  by  being  dissolved  in  it  (solution).  Similarly  a 
liquid  may  be  suspended  in  droplets  in  another  (emulsion),  as  in 
milk,  or  it  may  be  dissolved.  It  is  usually  easy  to  distinguish 
between  the  two  cases,  for  a  suspended  substance  settles  or  separates 
sooner  or  later,  while  a  dissolved  one  shows  no  such  tendency.  The 
cases  are  exceptional  where  the  subdivision  of  a  suspended  substance 
is  so  minute  (colloidal  solution)  as  to  make  its  retention  by  filter 
paper  impossible.  If  a  liquid  is  opalescent  or  opaque,  then  we  have 
a  case  of  suspension.  A  solution  is  a  clear,  transparent,  perfectly 
homogeneous  liquid,  in  which  the  dissolved  substance  seems  to  have 
been  dispersed  so  completely  that  the  liquid  cannot  be  distinguished 
by  the  eye  from  a  pure  substance. 

There  is  no  limit  to  the  amount  of  dissipation  which  may  thus 
be  produced.  A  single  fragment  of  potassium  permanganate,  for 
example,  which  gives  a  very  deep  purple  solution  in  water,  may  be 
dissolved  in  a  liter  or  even  in  twenty  liters  of  water,  and  the  purple 
tinge  which  it  gives  to  the  liquid  will  still  be  perfectly  perceptible  in 
every  part  of  the  larger  volume.  The  qualitative  characteristics, 
therefore,  of  solution  are  absence  of  settling,  homogeneity,  and  extremely 
minute  subdivision  of  the  dissolved  substance. 

The  Scope  of  the  Word.  —  The  word  is  used  for  other  systems 
than  those  containing  a  solid  body  dissolved  in  a  liquid.  Thus, 
liquids  also  may  be  dissolved  in  liquids,  as  alcohol  in  water.  Again, 


SOLUTION  97 

if  we  warm  ordinary  water,  bubbles  of  gas  appear  on  the  sides  of 
the  vessel  before  the  water  has  approached  the  boiling-point.  They 
are  found  to  be  gas  derived  from  the  air.  Agitation  of  any  gas  with 
water  results  in  the  solution  of  a  large  or  small  quantity  of  the  gas, 
and  heat  will  usually  drive  the  gas  out  again.  It  appears  therefore 
that  solids,  liquids,  and  gases  can  equally  form  solutions  in  liquids. 
The  absorption  of  hydrogen  by  palladium  (at  all  events  after  a 
certain  point),  and  by  iron,  takes  place  in  accordance  with  the  same 
laws  as  the  solution  of  solids  in  liquids,  and  the  results  may  be 
described  therefore  as  true  solutions.  Liquids  are  in  some  cases 
absorbed  by  solids,  and  homogeneous  mixtures  of  solids  with  solids 
are  perfectly  familiar.  The  sapphire  is  a  solution  of  a  small  amount 
of  a  strongly  colored  substance,  in  a  large  amount  of  colorless  alumi- 
nium oxide.  It  may  therefore  be  stated  that  solution  of  gases, 
liquids,  and  solids  in  solids  appears  to  be  possible. 

Limits  of  Solubility.  —  The  next  question  which  naturally  occurs 
to  us  is  as  to  whether  the  mingling  of  two  substances  in  this  manner 
has  any  limits.  We  find  that  the  results  of  experiment  in  this 
direction  may  be  divided  into  two  classes.  Some  pairs,  of  liquids 
particularly,  may  be  mixed  in  any  proportions  whatever.  '  Alcohol 
and  water  is  such  a  pair.  On  the  other  hand,  at  the  ordinary  labora- 
tory temperature,  we  can  scarcely  take  a  fragment  of  marble  (CaC03) 
so  small  that  it  will  dissolve  completely  in  100  c.c.  of  pure  water. 
Under  the  same  conditions  any  amount  of  potassium  chlorate  up 
to  5  g.  will  almost  completely  disappear  after  vigorous  stirring,  while 
90  g.  of  ordinary  Epsom  salts  (hydrated  magnesium  sulphate),  but 
not  more,  may  be  dissolved  in  about  the  same  amount  of  water.  In 
fact,  most  solids  may  be  dissolved  in  a  liquid  only  up  to  a  certain 
limit,  which  with  different  solids  may  range  from  a  scarcely  percep- 
tible to  a  very  large  amount.  No  substance  is  absolutely  insoluble. 
But  for  the  sake  of  brevity  we  call  marble,  for  example,  "  insoluble  " 
because  in  most  connections  it  may  be  so  considered. 

Recognition  and  Measurement  of  Solubility.  —  The  only 
method  of  recognizing  with  certainty  whether  a  solid  is  soluble  in  a 
liquid  or  not  is  to  filter  the  mixture  and  evaporate  a  few  drops  of 
the  filtrate  on  a  clean  watch-glass.  For  learning  how  much  of  the 
body  is  contained  in  a  given  solution,  a  weighed  quantity  of  the 


98  COLLEGE  CHEMISTRY 

solution  is  evaporated  to  dryness  and  the  weight  of  the  residue 
determined. 

It  must  be  stated  explicitly  that  in  going  into  solution,  as  we  have 
used  the  term,  a  compound  dissolves  as  a  whole  and,  if  the  compound 
is  pure  (p.  23),  any  residue  has  the  same  chemical  composition  as 
the  part  which  has  dissolved. 

Terminology.  —  In  order  to  describe  the  relations  of  the  compo- 
nents of  a  solution,  certain  conceptions  and  corresponding  technical 
expressions  are  required.  It  is  customary  to  speak  of  the  substance 
which,  like  water  in  most  cases,  forms  the  bulk  of  the  solution,  as  the 
solvent.  To  express  the  substance  which  is  dissolved,  the  word 
solute  is  frequently  used,  and  will  be  employed  when  we  wish  to 
avoid  circumlocution.  The  amount  of  the  substance  which  has  been 
dissolved  by  a  given  quantity  of  the  solvent  is  described  as  the  con- 
centration of  the  solution.  A  solution  containing  a  small  proportion 
of  the  dissolved  body  is  called  dilute;  it  has  a  small  concentration. 
One  which  contains  a  larger  amount  is  more  concentrated.  Very 
"  strong  "  solutions  are  frequently  spoken  of  simply  as  concentrated 
solutions.  The  partial  removal  of  the  solvent  (as  by  evaporation) 
is  called  concentrating,  its  total  removal  evaporating  to  dryness. 
Finally,  since  there  is  a  limit  to  the  solubility  of  most  substances, 
a  solution  is  described  as  saturated  when  the  solute  has  given  as 
much  material  to  the  solvent  as  it  can.  This  state  is  reached  after 
prolonged  agitation  with  an  excess  of  the  gas,  of  the  liquid,  or  of 
the  finely  powdered  solid,  as  the  case  may  be  (see  pp.  102,  107). 
Other  things  being  equal,  the  larger  the  excess,  the  sooner  satura- 
tion is  attained.  The  maximum  concentration  attainable  in  this 
way  is  called  the  solubility  of  the  substance  in  a  given  solvent. 

The  concentrations  of  solutions,  saturated  and  otherwise,  are  some- 
times expressed  in  physical,  and  sometimes  in  chemical,  units  of 
weight.  When  physical  units  are  employed,  we  give  the  number 
of  grams  of  the  solute  held  in  solution  by  one  hundred  of  the 
solvent. 

When  chemical  units  of  weight  are  employed,  two  different  plans 
are  possible,  and  both  are  in  use.  Either  the  equivalent  (p.  32)  or 
the  atomic  weights  may  be  taken  as  a  basis  of  measurement.  In  the 
former  case,  the  solutions  are  called  normal  solutions,  and  in  the  latter, 
for  a  reason  which  will  appear  later  (Chap,  xii),  molar  solutions. 


SOLUTION  99 

A  normal  solution  contains  one  gram-equivalent  of  the  solute  in  one 
liter  of  solution.  The  word  "  equivalent  "  has  been  used  hitherto 
only  of  elements,  and  this  application  of  the  expression  involves  an 
extension  of  its  meaning.  An  equivalent  weight  of  a  compound  is  that 
amount  of  it  which  will  interact  with  one  equivalent  of  an  element. 
Thus,  a  formula-weight  of  hydrochloric  acid  HC1  (36.5  g.)  is  also  an 
equivalent  weight,  for  it  contains  1  g.  of  hydrogen,  and  this  amount 
of  hydrogen  is  displaceable  by  one  equivalent  weight  of  a  metal.  A 
formula-weight  of  sulphuric  acid  H2SO4  (98  g.),  however,  contains 
two  equivalents  of  the  compound,  and  a  formula-weight  of  alumi- 
nium chloride  A1C13  (133.5  g.)  three  equivalents.  Hence  normal 
solutions  of  these  three  substances  contain  respectively  36.5  g.  (HC1), 
49  g.  (H2SO4),  and  44.5  g.  (A1C13)  per  liter  of  solution.  The  special 
property  of  normal  solutions  is,  obviously,  that  equal  volumes  of  two 
of  them  contain  the  exact  proportions  of  the  solutes  which  are 
required  for  complete  interaction.  Solutions  of  this  kind  are  much 
used  in  quantitative  analysis.  We  frequently  use  also  decinormal 
or  one-tenth  normal  solutions  (.IN  or  A7"/ 10),  and  seminormal 
(.5  N  or  2V/ 2),  and  six  times  normal  solutions  (6  N),  and  so  forth. 

A  molar  solution  contains  one  mole  (gram-molecular  weight)  of  the 
solute  in  one  liter  of  solution.  When  molecular  formulae  (see  Chap, 
xii)  are  used,  this  means  one  gram-formula  weight  per  liter.  In  the 
cases  cited  above,  the  molar  solution  contains  36.5  g.  (HC1),  98  g. 
(H2S04),  and  133.5  g.  (A1C13)  per  liter.  As  will  be  seen,  the  concen- 
trations of  molar  and  normal  solutions  are  necessarily  identical  when 
the  radicals  are  univalent. 

The  solubilities  at  18°  of  one  hundred  and  forty-two  bases  and  salts 
are  given  in  a  table  printed  inside  the  cover,  at  the  front  of  this  book. 

Solution  One  of  the  Physical  States  of  Aggregation  of 
Matter.  —  When  a  solid  body  dissolves  in  a  liquid,  the  properties  of 
the  body  undergo  a  very  marked  change,  which  to  all  appearance 
might  be  chemical.  Yet,  besides  the  ease  with  which  a  liquid  may  be 
removed  by  evaporation  and  the  solid  recovered  unchanged,  we  note 
particularly  that  the  concentration  of  a  saturated  solution  cannot  be 
expressed  in  terms  of  integral  multiples  of  the  chemical  combining 
weights.  We  shall  see  also  that  the  quantity  of  a  solid  which  a 
liquid  may  take  up  varies  with  the  slightest  changeJn  temperature. 
Now  we  do  not  find  the  composition  <tf^ckQmi$aE  compounds  »so  Jtb 


100  COLLEGE    CHEMISTRY 

vary.  The  solution  of  a  solid  may  therefore  be  likened  to  a  change 
in  state  of  aggregation,  similar  to  the  conversion  of  a  liquid  into  a 
gas  or  a  solid. 

As  in  other  changes  of  state,  so  in  the  process  of  solution,  heat  is 
always  liberated  or  absorbed.  This  is  known  as  heat  of  solution. 
Thus,  one  formula-weight  of  sulphuric  acid,  in  dissolving  in  a  large 
volume  of  water,  liberates  39,170  calories,  and  one  formula-weight 
of  ammonium  chloride,  in  dissolving,  absorbs  3880  calories. 

As  there  is  danger  of  confusion  arising,  we  may  repeat  that  a  com- 
pound is  homogeneous  and  its  composition  is  expressible  in  chemical 
units  of  weight;  a  saturated  solution  is  homogeneous  but  its  concen- 
tration varies  with  temperature  so  that  chemical  units  cannot  be 
used  to  describe  its  composition;  a  mixture  of  two  solids,  or  an  emul- 
sion of  two  liquids,  is  neither  homogeneous  nor  in  any  way  definite 
in  composition. 


Kinetic-Molecular  Hypothesis  Applied  to  the  State  of  Solu- 

Accepting  solution  as  a  physical  state  of  aggregation,  we 
may  now  apply  the  same  formulative  hypothesis  to  the  explanation 
of  the  behavior  of  a  substance  in  solution  as  to  matter  in  the  gaseous 
or  liquid  states.  We  saw  that  a  solid  body,  which  is  ordinarily  con- 
densed in  a  small  space,  can  be  disseminated  by  the  use  of  a  solvent 
through  a  very  large  one.  The  molecules  of  the  solid  become  scat- 
tered like  those  of  a  gas  or  vapor  through  a  much  greater  volume. 
We  may  regard  the  dissolved  substance  as  being,  practically,  in  a 
gaseous  or  quasi-gaseous  condition.  The  molecules  are  torn  apart 
from  one  another,  their  cohesion  is  overcome,  and  their  freedom  of 
motion  is  in  a  measure  restored.  It  is  true  that  they  could  not 
continue  to  occupy  this  large  volume  for  a  moment  in  the  absence  of 
the  solvent.  But  we  may  bring  this  into  relation  with  the  case  of  a 
vapor  by  saying  that  a  solid  body,  like  common  salt,  can  only  evapo- 
rate (i.e.  "dissolve")  at  the  ordinary  temperature,  and  occupy  a  large 
space,  when  that  space  is  already  filled  with  a  suitable  liquid.  The 
latter  acts  as  a  vehicle  for  the  particles  of  the  solid.  A  volatile 
liquid,  on  the  contrary,  can  dissolve  in  an  empty  space  and  fill  it  with 
its  particles  without  any  vehicle  being  required. 

This  conception  of  the  quasi-gaseous  condition  of  a  dissolved  sub- 
stance woujd  be^  simply  fantastic  if  it  did  not  lead  us  to  a  better 
g  of  ;'tHe  betffraoi1  !o£  solutions.     It  does  successfully 


SOLUTION 


101 


explain  many  things,  such  as  diffusion,  osmotic  pressure,  and  satura- 
tion (see  next  section). 

It  is  easy  to  show  that,  if  we  place  a  quantity  of  the  pure  solvent 
(Fig.  37)  above  a  concentrated  solution  of  a  substance,  and  then  set 
the  arrangement  aside,  the  dissolved  body  slowly  makes  its  way 
through  the  liquid  (Fig.  38),  obliterating  the  original  plane  of  sepa- 
ration. Eventually  the  dissolved  body  scatters  itself  uniformly 
through  the  whole.  In  other  words,  the  particles  of  the  dissolved 


FIG.  38. 


substance  exhibit  the  .property  of  diffusion  in  the  same  way  as  do 
those  of  gases. 

When  the  diffusion; of  a  gas  is  resisted  by  a  suitable  partition,  we 
find  pressure  is  exercised  upon  the  walls  of  the  vessel  and  upon  the 
partition.  It  is  possible  to  show  that  the  particles  of  a  dissolved 
substance  exercise  a  pressure  of  a  very  similar  kind.  This  pressure 
is  spoken  of  as  osmotic  pressure.  This  pressure  is  found  to  be  pro- 
portional to  the  concentration  of  the  solution,  just  as  gaseous 
pressure  is  proportional  to  the  concentration  of  the  gas  (Boyle's  law). 
y 

Kinetic- Molecular  Hypothesis  Applied  to  the  Process  of 
Solution.  —  We  may  now  apply  the  same  hypothesis  to  the  process 
of  dissolving,  with  a  view  more  especially  to  explaining  why  the 


102 


COLLEGE   CHEMISTRY 


process  of  dissolving  ceases,  in  spite  of  the  presence  of  excess  of  the 
solute,  when  a  certain  concentration  has  been  reached.  If  some  of 
the  material  dissolves,  why  not  more? 

Let  us  suppose  that  it  is  the  dissolving  of  common  salt  in  water 
(Fig.  39)  which  we  wish  to  explain  in  detail.     We  believe  that  in 
the  solid  substance  the  molecules  are  somewhat  closely  packed 
together,  while  in  the  solution  they  are  rather 
sparsely  distributed.     The  process  of  solution 
must  consist  in  the  loosening  of  the  molecules 
on  the  surface  and  their  passage  into  the  liquid. 
By  diffusion,  the  free  molecules  will  gradually 
move  away  from  the  neighborhood  of  the  sur- 
face of  the  solid  and  make  room  for  others, 
and  thus,  if  the  system  remains  undisturbed, 
the  liquid  will  eventually  become  a  solution 
of  uniform  concentration.     If  a  large  enough 
amount  of  the  solid  has  been  provided,  the 
ultimate  condition  will  be  that  of  a  saturated 
solution  with  excess  of  the  solid  beneath.     If 
we  had  proper   means   of   measuring  it,   the 
/  tendency  of  the  molecules  to  leave  the  solid  in 
the  presence  of  a  given  liquid  would  give  the 
FIG.  39.         r      effect  of  a  kind  of  pressure.     This  is  spoken  of 

J  as  solution  pressure. 

Now  the  molecules,  after  having  entered  the  liquid,  move  in  every 
direction,  and  consequently  some  of  them  will  return  to  the  solid 
and  attach  themselves  to  it.  The  frequency  with  which  this  will 
occur  will  be  greater  as  the  crowding  of  particles  in  the  liquid 
increases,  so  that  a  stage  will  eventually  be  reached  at  which  the 
number  of  molecules  leaving  the  solid  will  be  no  greater  than  that 
landing  upon  it  in  a  given  time.  If  the  whole  of  the  liquid  has 
meanwhile  become  equally  charged  with  dissolved  molecules,  there 
will  be  no  chance  that  the  field  of  liquid  immediately  round  the  solid 
will  lose  them  by  diffusion,  so  that  a  condition  of  balance  or  equili- 
brium (p.  91)  will  have  been  established:  NaCl  (solid)  <=±  NaCl 
(dslvd).  The  motion  of  the  particles  in  the  liquid  produces  what  we 
have  called  osmotic  pressure;  and  when  the  osmotic  pressure,  by  the 
continual  increase  in  the  number  of  dissolved  molecules,  becomes 
equal  to  the  solution  pressure,  increase  in  concentration  of  the  solution 


SOLUTION  103 

ceases.  It  is  at  this  point  that  we  speak  of  the  solution  as  being  satu- 
rated with  respect  to  the  particular  substance  dissolving.  The 
analogy  to  vapor  tension  and  vapor  pressure  (p.  79)  is  evident. 
The  foregoing  explanation  should  be  compared  carefully  with 
that  given  in  the  section  on  the  kinetic  hypothesis  applied  to 
liquids,  and  in  that  on  equilibrium  (pp.  90-93). 

When  the  dissolving  substance  is  a  gas,  led  through,  or  confined 
above  the  liquid  at  a  definite  pressure,  the  gas  dissolves  until  a  state 
)f  equilibrium  between  dissolving  and  emission  is  reached,  for 
example,  Oxygen  (gas)  <=±  Oxygen  (dslvd),  and  the  liquid  is  then 
saturated  with  the  gas. 

It  is  found,  as  the  hypothesis  would  lead  us  to  expect,  that  the 
concentration  of  the  saturated  solution  of  a  gas  is  proportional  to  the 
pressure  at  which  the  gas  is  supplied  (Henry's  law). 

Independent  Solubility.  —  Just  as  two  gases,  when  mixed,  are 
independent  of  one  another  (p.  60),  and  have  severally  the  same 
pressure,  solubility,  and  so  forth,  as  they  would  possess  if  each  alone 
occupied  the  same  space,  so  is  it  with  dissolved  substances.  In  gen- 
eral, a  volume  of  water,  in  which  a  moderate  amount  of  some  sub- 
stance has  been  dissolved,  will  take  up  as  much  of  a  second  substance 
as  would  an  equal  volume  of  pure  water.  Thus,  water  containing 
some  sugar  will  dissolve  as  much  sodium  chloride  as  the  same  amount 
of  pure  water.  In  the  point  of  view  of  the  kinetic-molecular  hypoth- 
esis, the  dissolved  molecules  of  sugar  have  no  connection  with,  or 
influence  upon,  the  mechanism  which  determines  the  solubility  of 
the  salt,  namely,  the  exchange  of  salt  molecules  between  the  sus- 
pended, dissolving  crystals  and  the  solution. 

Two  Immiscible  Solvents :  Law  of  1*art(tion. — An  interesting 
application  of  the  same  ideas  may  be  made  to  a  case  which  occurs 
very  commonly  in  chemical  work.  If  we  shake  up  a  small  particle 
of  iodine  with  water,  we  find  that  it  dissolves  slowly,  giving  eventu- 
ally a  saturated  but  very  dilute  solution.  If  now  ether  in  sufficient 
quantity  be  shaken  with  the  aqueous  solution,  the  greater  part  of 
the  iodine  will  find  its  way  into  the  ether,  and  be  contained  in  the 
brown  layer  which  rises  to  the  top.  The  process  of  removing  a 
substance  partially  from  solution  in  one  solvent  and  securing  it  in 
another  is  called  extraction.  We  find  in  such  cases  that  neither  sol- 


104 


COLLEGE    CHEMISTRY 


90°  100 e 


SOLUTION  105 

| 

vent  can  entirely  deprive  the  other  of  the  whole  of  the  dissolved 
substance,  if  the  latter  is  soluble  in  both  independently.  A  state  of 
equilibrium  is  finally  reached:  I  (in  Aq)  <=±  I  (in  ether).  The  parti- 
tion of  the  substance  takes  place  in  proportion  to  its  solubility  in 
each  solvent.  It  is  found  that  any  amount  of  the  solute,  up  to  the 
maximum  the  system  can  contain,  provided  this  does  not  involve 
too  high  a  concentration  in  either  solvent,  is  divided  so  that  the  ratio 
of  the  concentrations  in  the  two  solvents  is  always  the  same.  In  the 
case  of  iodine  divided  between  water  and  ether,  this  ratio  is  about 
1  :  200. 

Influence  of  Temperature  on  Solubility.  —  The  quantity  of  a 
substance  which  we  can  dissolve  in  a  fixed  amount  of  a  given  solvent 
depends  very  largely  upon  the  temperature  of  both.  Usually  the 
solubility  increases  with  rise  in  temperature.  Measurements  may 
be  made  by  the  method  described  before  (p.  97),  using  excess  of  the 
finely  powdered  solute  with  different  portions  of  the  same  solvent 
in  vessels  kept  at  different  temperatures.  The  most  useful  way  of 
representing  the  results  is  to  plot  them  graphically.  The  diagram 
(Fig.  40)  shows  the  curves  for  a  few  familiar  substances.  The  ordi- 
nates  represent  the  number  of  grams  of  the  anhydrous  compound 
which  is  held  in  solution  by  100  g.  of  water  in  each  case.  The 
abscissae  represent  the  temperatures.  The  concentration  for  any 
temperature  can  be  read  off  at  once.  Thus  100  g.  of  water  holds 
13  g.  of  potassium  nitrate  in  solution  at  0°  and  150  g.  at  73°.  The 
increase  in  solubility  is  here  enormous.  On  the  other  hand,  the 
same  quantity  of  water  will  hold  35.6  g.  of  sodium  chloride  in  solu- 
tion at  0°  and  39  g.  at  100°.  The  difference  is  shown  at  once  when 
we  examine  the  curves  and  observe  that  the  line  representing  the 
solubility  of  sodium  chloride  scarcely  rises  at  all  between  0°  and 
100°,  while  that  of  potassium  nitrate  is  extremely  steep. 

Cases  in  which  the  solubility  decreases  with  rise  in  temperature  are 
less  common.  Anhydrous  sodium  sulphate  (Fig.  41,  p.  106)  is  an 
illustration. 

Equilibrium  in  a  Saturated  Solution.  —  Once  a  solution  has 
become  saturated,  the  dissolving  substance  remains  thereafter  un- 
changed in  amount.  A  greater  excess  of  the  solute  forces  no  more 
matter  into  solution  than  does  a  small  excess. 


106 


COLLEGE   CHEMISTRY 


It  should  be  clearly  understood,  however,  that  the  kinetic  hypothe- 
sis requires  us  to  assume  that  an  exchange  of  molecules  (p.  102)  is 
still  going  on  between  the  solid  and  the  solution.  That  this  concep- 
tion is  correct  may  be  shown  in  vari&us  ways.  Thus,  if  a  crystal, 
the  edges  or  corners  of  which  have  been  broken,  is  suspended  in  a 
saturated  solution  of  the  same  substance,  it  neither  increases  nor 
diminishes  in  weight.  Yet  we  find  that  the  imperfections  are 
removed,  and  that  this  takes  place  by  the  solution  of  a  portion  of 
the  substance  from  the  perfect  surfaces  and  its  deposition  upon  the 
imperfect  ones. 

Another  very  striking  proof  of  this  may  be  obtained  by  saturating 
water  with  ordinary  Glauber's  salt  (hydrated  sodium  sulphate, 


00      10°     20°      30° 


50°      60°      70°      80°     90°    100° 
Temperature 

FIG.  41. 


10H2O)  at  a  temperature  somewhat  above  the  ordinary, 
say  30°.  The  excess  of  the  solid  is  carefully  and  completely  sepa- 
rated from  the  liquid,  and  the  latter  is  allowed  to  cool  in  a  flask 
loosely  stoppered  with  cotton.  The  solution  now  contains  (Fig.  41) 
a  much  larger  amount  of  sodium  sulphate  (Na-^SOJ  than  at  its 
present  temperature  it  could  acquire  from  contact  with  Glauber's 
salt.  Yet  in  the  absence  of  a  crystal,  with  which  the  above  described 
exchange  could  take  place,  no  deposition  of  the  dissolved  substance 
begins.  The  solution  may  be  kept  indefinitely  without  alteration. 
The  introduction,  however,  of  the  minutest  fragment  of  the  deca- 


SOLUTION  107 

hydrate  at  once  starts  the  exchange,  and  this  is  necessarily  very 
much  to  the  disadvantage  of  the  solution  and  the  advantage  of  the 
crystal:  Na2SO4  (dslvd)  <=±  Na2SO4,10H2O  (solid).  The  latter  there- 
fore forms  the  center  of  a  radiating  mass  of  blade-like  processes, 
which  sprout  with  astonishing  rapidity  through  the  liquid. 

Usually  the  cooling  of  a  concentrated  solution  leads  to  the  almost 
immediate  appearance  of  crystals  spontaneously,  and  the  substance 
is  deposited  gradually  as  the  temperature  falls.  But  solutions  of  a 
number  of  common  substances,  such  as  sodium  thiosulphate  (photo- 
grapher's "  hypo  ")  and  sodium  chlorate,  behave  like  that  of  sodium 
sulphate.  They  are  said  to  have  a  tendency  to  give  supersaturated 
solutions.  In  general,  crystallization  can  be  started  only  by  intro- 
duction of  a  specimen  of  the  same  substance.  The  smallest  particle 
of  the  right  material  floating  in  the  air,  if  it  gains  admission,  will 
bring  about  the  result.  This  shows  the  importance  of  the  interchange 
of  molecules,  of  which  we  have  spoken,  for  establishing  equilibrium. 

To  avoid  a  common  misconception,  it  must  be  noted  that  a  saturated 
solution  is  not  one  containing  all  of  the  solute  that  it  can  hold.  A 
supersaturated  solution  contains  more.  The  saturated  solution  is  one 
which  contains  all  of  the  dissolved  solute  that  it  can  acquire  from 
the  undissolved  solute.  Better  still,  it  is  that  solution  which,  when 
placed  with  excess  of  the  solute,  is  found  to  be  in  equilibrium. 

Exercises. —  1.  Give  other  examples  of  limited  solubility  in 
various  solvents  (p.  97). 

2.  What  weights  of  phosphoric  acid  (p.  51)  and  of  sodium  hydr- 
<  oxide,  respectively,  are  required  to  make  1  liter  of  a  normal  solution? 

3.  Express  the  concentrations  of  solutions  of  ammonium  chloride, 
saturated  at  0°  (sp.  gr.  1.076),  and  of  potassium  sulphate  K2SO,, 
saturated  at  10°  (sp.  gr.  1.083),  in  terms  of  a  normal  solution  (p.  99). 

4.  Express  the  concentration  of  a  five  per  cent  aqueous  solution 
of  phosphoric  acid  (sp.  gr.  1.027),  in  terms  of  a  normal  and  a  molar 
solution,  respectively. 

5.  Explain  why,  (a)  pulverization  and,  (6)  agitation  hasten  the 
dissolving  of  a  solid? 

6.  Read  from  the  curves  (p.  104)  the  solubilities  of  potassium 
nitrate  at  15°,  of  potassium  chloride  at  30°,  of  potassium  chlorate  at 
45°.     What  are  the  relative  rates  at  which  the  solubilities  of  these 
salts  increase  with  rise  in  temperature? 

lid 


V. 


CHAPTER  XI 
CHLORINE   AND   HYDROGEN    CHLORIDE 

CHLORINE  was  first  recognized  as  a  distinct  substance  by  Scheele 
(1774).  He  obtained  it  from  salt  by  means  of  manganese  dioxide, 
using  the  method  described  below. 

Occurrence.  —  Chlorine  does  not  occur  free  in  nature.  There  are, 
however,  many  compounds  of  it  to  be  found  in  the  mineral  kingdom. 
Sea-water  contains  a  number  of  chlorides  in  solution.  Of  the  3.6 
per  cent  of  solid  matter  in  sea-water,  nearly  2.8  is  common  salt 
(sodium  chloride,  NaCl).  During  past  geological  ages  the  evapora- 
tion of  sea-water  has  led  to  the  formation  of  immense  deposits  of 
the  compounds  usually  found  in  such  water.  Thus,  at  Stassfurt, 
such  strata  attain  a  thickness  of  over  a  thousand  feet.  Certain 
layers  of  these  strata  are  composed  mainly  of  sodium  chloride,  called 
by  the  mineralogist  halite  (rock  salt).  In  other  layers  potassium 
chloride  (sylvite),  and  hydrated  magnesium  chloride  (bischofite), 
and  other  compounds  of  chlorine,  occur.  The  chloride  of  silver  (horn 
silver)  is  a  valuable  ore. 

Preparation.  —  Chlorine  cannot  be  obtained  with  the  same  ease 
as  oxygen.  There  are  only  a  few  chlorides,  such  as  those  of  gold  and 
platinum,  which  lose  chlorine  when  heated,  and  they  are  too  expen- 
sive or  difficult  to  make  for  laboratory  use.  We  employ  therefore 
methods  like  those  used  for  the  preparation  of  hydrogen  (cf.  p.  67). 
We  may  (1)  decompose  any  chloride  by  means  of  electricity,  just  as, 
to  get  hydrogen,  we  electrolyzed  a  dilute  acid  (p.  64).  Or  (2)  we 
may  take  some  inexpensive  compound  of  chlorine,  such  as  hydrogen 
chloride  (HC1),  and  by  means  of  some  simple  substance  which  is 
capable  of  uniting  with  the  other  constituent,  —  here  oxygen  serves 
the  purpose,  —  secure  the  liberation  of  the  element.  Or  (3)  —  and 
this  turns  out  to  be  the  most  convenient  laboratory  method  —  we 
may  use  a  more  complex  action. 

108 


CHLORINE 


109 


Electrolysis  of  Chlorides.  —  Hydrogen  chloride  and  those  chlo- 
rides of  metals  which  are  soluble  in  water  are  all  decomposed  when 
a  current  of  electricity  is  passed  through  the  aqueous  solution.  They 
yield  chlorine  at  the  positive  electrode.  The  other  constituent,  the 
hydrogen  (Fig.  42),  manganese,  or  whatever  it  may  be,  is  liberated 
at  the  negative  wire.  To  decompose  hydrochloric  acid  an  electro- 
motive force  of  at  least  1.31  volts  is  required.  Since  the  chlorine  is 
soluble  in  water,  the  effervescence  due  to  its  release  is  not  noticeable 
until  the  liquid  round  the  electrode  has  become  saturated  with  the 
gas:  C12  (dslvd)  <=±  C12  (gas).  The  shape  of  the  apparatus  keeps  the 
two  products  from  mingling. 
The  presence  of  the  chlorine 
in  the  liquid  at  the  positive  end 
may  be  shown  by  a  suitable  test 
(pp.  66  and  114). 

In  commerce  chlorine  is  now 
obtained  chiefly  by  this  method, 
sodium  chloride  or  potassium 
chloride  being  the  source  of  the 
element.  Electrodes  of  arti- 
ficial graphite  are  used,  as  most 
other  conductors  unite  with  the 
chlorine.  The  potassium  or 
sodium,  as  the  case  may  be, 
travels  towards  the  negative 
electrode,  but  is  not  liberated.  FIG.  42. 

Instead,  potassium  or  sodium 

hydroxide  accumulates  in  the  solution  round  the  plate  and  hydro- 
gen escapes.  The  chlorine  is  released  at  the  positive  electrode,  as 
usual.  The  hydroxide  and  the  chlorine  both  find  chemical  appli- 
cations. The  chlorine  is  either  liquefied  by  compression  in  steel 
cylinders  or  employed  at  once  for  making  bleaching  powder  (q.v.). 

Action  of  Free  Oxygen  on  Chlorides.  —  Sodium  chloride  is  the 
cheapest  source  of  chlorine,  but  oxygen  does  not  interact  with  it 
even  at  a  high  temperature.  By  treating  the  sodium  chloride  with 
sulphuric  acid,  therefore,  the  chlorine  is  first  transferred  into  com- 
bination with  the  hydrogen  of  the  acid.  The  details  of  this  action 
are  described  below  (see  p.  117).  In  order  to  liberate  chlorine  from 


110 


COLLEGE  CHEMISTRY 


the  hydrogen  chloride,  we  may  then  combine  the  hydrogen  with 
oxygen  obtained  from  the  air.  The  action  is  in  accordance  with  the 
equation : 

2HC1  +  O  «=>  H2O  +  2C1. 

The  two  gases  interact  so  slowly,  however,  that  a  catalytic  agent 

must  be  employed.  The 
mixture  of  air  and  hydro- 
gen chloride  is  passed 
over  pieces  of  heated 
pumice-stone  or  broken 
brick  previously  saturated 
with  cupric  chloride  solu- 
tion. A  temperature  of 
370°-400°  is  used.  Fur- 
thermore, the  action  is 
reversible  and  equilibrium 
is  reached  when  80  per 
cent  of  the  hydrogen 
chloride  has  been  decom- 
FIG.  43.  posed.  Hence  20  per  cent 

of    this    gas    passes    on 

unchanged.  In  the  product,  the  chlorine  is  mixed  with  steam  and 
with  a  very  large  volume  of  nitrogen  which  entered  with  the 
oxygen,  as  well  as  with  unused  hydrogen  chloride,  so  that,  for 
making  the  pure  substance,  this  method  (Deacon's  process)  is  quite 
unsuitable. 

The  above  action  is  spoken  of  as  an  oxidation.  It  is  true  that  no 
oxygen  is  actually  introduced  into  the  hydrogen  chloride  as  a  whole. 
The  removal  of  hydrogen  from  combination  with  the  chlorine  is, 
however,  the  first  step  towards  the  introduction  of  oxygen  into  com- 
bination with  the  latter,  and  is  essentially  an  oxidation. 

Action  of  Combined  Oxygen  upon  Chlorides.  —  The  best 
laboratory  method  for  making  chlorine  is  to  place  some  solid 
potassium  permanganate  in  a  flask,  arranged  like  that  in  Fig.  43. 
Concentrated  commercial  hydrochloric  acid  (an  aqueous  solution 
of  hydrogen  chloride),  diluted  with  one-third  of  its  volume  of 
water,  is  allowed  to  fall  upon  the  compound  drop  by  drop  from  the 


CHLORINE  111 

dropping  funnel.  The  action  is  very  rapid,  the  acid  is  exhausted 
almost  as  fast  as  it  falls,  and  so  the  stream  of  gas  can  be  stopped 
by  simply  closing  the  stopcock.  The  gas  is  passed  through  a 
washing  bottle  containing  water,  in  order  to  remove  some  hydro- 
gen chloride  which  may  be  carried  over.  It  may  be  dried,  if 
necessary,  in  a  second  washing  bottle  containing  concentrated 
sulphuric  acid.  It  cannot  be  collected  over  water  on  account  of 
its  solubility,  so  that  jars  are  usually  filled  with  it  by  downward 
displacement  of  air. 

The  skeleton  equation  (p.  42)  here  is: 

KMnO4  +  HC1  ->  H20  +  KC1  +  MnCl2  +  Cl. 

The  O4,  being  all  converted  into  water,  requires  8H,  and  therefore 
8HC1,  for  the  action.  The  two  metals,  potassium  and  manganese, 
give  their  respective  chlorides,  KC1  and  MnCl2.  This  uses  3C1,  and 
hence  5C1  remains  over  to  be  liberated: 

vxKMnO4  +  8HC1  ->  4H2O  +  KC1  +  MnCl2  +  5C1. 

The  combined  oxygen  of  the  permanganate  has  oxidized  the  hydro- 
gen chloride,  just  as  did  the  free  oxygen  in  Deacon's  process. 

Many  other  compounds  of  oxygen  with  metals  interact  with  hydro- 
chloric acid  to  give  free  chlorine.  Lead  dioxide  PbO2,  potassium 
chlorate  KC1O3,  potassium  dichromate  K2Cr2O7,  and  manganese 
dioxide  MnO2,  are  of  this  nature.  The  last,  being  inexpensive,  is 
commonly  used  in  making  chlorine.  Being  an  insoluble  substance, 
however,  the  manganese  dioxide  acts  much  more  slowly  than  does 
the  potassium  permanganate,  which  is  soluble.  A  large  amount  of 
the  materials,  and  the  aid  of  heat,  are  required  to  secure  a  rapid 
stream  of  chlorine. 

The  action  of  manganese  dioxide  upon  hydrochloric  acid  is  an 
instructive  one.  It  is  a  general  rule,  of  which  we  shall  meet  many 
applications,  that  when  an  acid  interacts  with  an  oxide  of  a  metal,  there 
are  two  constant  features  in  the  result,  namely:  (1)  The  oxygen  of 
the  oxide  combines  with  the  hydrogen  of  the  acid  to  form  water,  and  (2) 
the  metal  of  the  oxide  combines  with  the  acid  radical  of  the  acid  accord- 
ing to  the  valences  of  each.  Here  the  skeleton  equation  should  be, 
Mn02  +  HC1  ->  H20  +  MnCl4.  For  O2,  to  form  water,  4HC1  is 
required,  and  the  product  is  2H2O.  Hence  the  equation  might  be: 

Mn02  -f  4HC1  ->  2H20  +  MnCl4. 


\ 


112  COLLEGE   CHEMISTRY 

This  is  what  happens  in  the  first  place.  The  products  actually 
obtained,  however,  are  water,  manganous  chloride  (MnCl2)  and 
chlorine.  The  manganese  tetrachloride  is  decomposed  by  the  heat- 
ing, the  chlorine  escapes,  and  the  other  two  products  remain  in  the 
vessel. 

MnO2  +  4HC1  ->  2H2O  +  MnCl2  +  2C1.  (1) 

We  owe  the  chlorine  to  the  fact  that  the  tetrachloride  is  unstable. 

If  we  had  used  manganous  oxide  (MnO),  we  should  have  had  a 
double  decomposition: 

MnO  +  2HC1  ->  H20  +  MnCl2,  (2) 

but  we  should  have  got  no  chlorine.  Perhaps  the  simplest  way  to 
describe  the  difference  between  these  two  actions  is  in  terms  of  the 
valence  of  the  manganese.  In  MnIV  O2n  the  element  is  quadrivalent. 
This  means  that  its  atomic  weight  professes  to  be  able  to  hold  four 
unit  weights  of  a  univalent  element.  The  four  valences  of  oxygen 
(2On)  can  do  the  same  thing.  In  equation  (1)  the  oxygen  fulfills 
this  promise  by  taking  4H1.  But  the  MnIV  can  hold  only  2C11  per- 
manently and  lets  the  other  2C11  go  free.  In  other  words,  the 
valence  of  the  unit  weight  of  manganese  changes  in  the  course  of  the 
action.  In  equation  (2),  on  the  other  hand,  the  manganese  is  biva- 
lent to  start  with  (MnuOn),  and  is  able  to  retain  the  amount  of 
chlorine  (2CP)  equivalent  to  O".  Actions  like  that  of  manganese 
dioxide  in  (1)  are  classed  as  oxidations.  The  hydrogen  chloride,  or 
rather  half  of  it,  is  oxidized.  A  graphic  mode  of  writing  may  make 
this  remark  clearer: 

. ,  IV  //  O  +  2HC1  -»  H20  +  Mn"a 
^  O  +  2HC1  -»  H20  +  2C1. 

The  upper  half  is  a  double  decomposition,  the  lower  an  oxidation 
by  half  the  combined  oxygen  of  the  dioxide.  The  same  explana- 
tion applies  to  the  interaction  of  lead  dioxide  with  hydrochloric 
acid. 

In  practice,  instead  of  employing  aqueous  hydrochloric  acid  we 
frequently  use  the  materials  from  which  it  is  prepared,  namely, 
common  salt  and  concentrated  sulphuric  acid  (p.  117),  along  with 
the  manganese  dioxide.  Under  those  circumstances,  the  action 


CHLORINE;  11 3 

appears  more  complex,  but  is  simply  a  combination  of  the  two  chem- 
ical changes,  and  is  represented  by  the  equation: 

MnO2  +  2NaCl  +  3H2SO4  -» 2H2O  +  2NaHSO4  +  MnSO4  +  2C1. 

This  mixture  cannot  be  used  with  potassium  permanganate,  as 
explosion  is  apt  to  occur. 

Physical  Properties.  —  Chlorine  differs  from  the  gases  we  have 
encountered  so  far  in  having  a  strong  greenish-yellow  tint  (Gk. 
xAwpo's,  pale  green),  a  fact  which  gave  rise  to  its  name,  and  having  a 
powerful,  irritating  effect  upon  the  membranes  of  the  nose  and  throat. 

Density  (H  =  1),  35.79  Boiling-point  (liq.),  -  33.6° 

Weight  of  1  1.,  3.220  g.  Melting-point  (solid),  -  102° 

Sol'ty  in  Aq  (20°),  215  vols.  in  100  Vap.  tension  (liq.)  0°,  3.66  atmos. 

Grit,  temp.,  4-  146°  Vap.  tension  (liq.)  20°,  6.62  atmos. 

Since  a  liter  of  air  weighs  1.293  g.,  chlorine  is  two  and  a  half  times 
heavier.  In  solubility  it  stands  between  slightly  soluble  gases,  like 
oxygen  and  hydrogen,  and  those  which  are  extremely  soluble.  It 
can  be  collected  over  hot  water  or  a  strong  solution  of  salt. 

Chlorine  was  first  liquefied  by  Northmore  (1805).  It  forms  a 
yellow  liquid  which,  contained  in  steel  cylinders  lined  with  lead, 
is  now  an  article  of  commerce.  On  being  cooled  below  -  102°,  it 
gives  a  pale-yellow  solid. 

Chemical  Properties.—  Chlorine  is  at  least  as  active  a  sub- 
stance as  is  oxygen.  It  presents  a  more  varied  array  of  chemical 
properties  than  does  that  element. 

1.  Chlorine  unites  directly  with  many  elements.  A  jet  of  hydrogen 
.burns  vigorously  in  chlorine,  producing  hydrogen  chloride.  The 
union  of  the  gases,  when  a  mixture  of  them  is  kept  cold  and  in  the 
dark,  is  too  slow  to  be  perceived.  On  exposure  to  diffused  light, 
however,  they  unite  slowly,  while  a  sudden  flash  of  sunlight  or  the 
burning  of  a  magnesium  ribbon  causes  instant  explosion.  The  func- 
tion of  the  light  here  is  entirely  different  from  that  in  the  decompo- 
sition of  silver  chloride  (pp.  9,  13).  In  the  latter  case  light  was  used 
to  maintain  the  change,  which  comes  to  a  stop  whenever  the  light 
is  withdrawn.  The  action  was  endothermal  and  consumed  energy. 
The  union  of  hydrogen  and  chlorine  is  highly  exothermal,  and  a 
minimum  of  light  only  is  needed  to  start  it  (pp.  54,  55). 


114  COLLEGE   CHEMISTRY 

Sodium  burns  in  chlorine,  producing  a  cloud  of  white  particles  of 
sodium  chloride.  Copper  in  the  condition  of  thin  leaf,  commonly 
used  for  gilding  (Dutch-metal),  catches  fire  spontaneously  when 
thrust  into  the  gas.  Phosphorus  burns  in  it  with  a  rather  feeble 
light,  producing  primarily  phosphorus  trichloride  PC13,  a  liquid 
(b.-p.  74°).  If  excess  of  chlorine  is  present,  then,  as  the  trichloride 
cools,  it  combines  to  form  the  solid  pentachloride  PC15.  Almost  all 
the  more  familiar  elements  unite  directly  with  chlorine  to  form 
chlorides.  The  exceptions  are,  nitrogen,  oxygen,  carbon,  and  the 
helium  group  of  elements  (p.  49).  Compounds  with  all  but  the 
last  may  be  obtained  as  products  of  more  complex  interactions,  how- 
ever. Some  elements,  like  phosphorus,  form  more  than  one  chloride. 

Carefully  dried  substances,  even  when  heated,  unite  slowly  or,  to 
all  appearance,  not  at  all  with  dry  chlorine.  The  introduction  of  a 
drop  of  water,  however,  into  a  remote  part  of  the  apparatus  at  once 
supplies  the  trace  of  moisture  which  seems  to  be  necessary  to  facili- 
tate the  chemical  change.  The  water  is  a  catalytic  agent,  and  we 
regard  it  as  simply  hastening  an  action  which  otherwise  is  vauishingly 
slow  (p.  54). 

2.  Chlorine,  like  sodium  (p.  66),  may  also  displace  elements  which 
are  already  in  combination.  Thus,  when  turpentine  (CioHi6)  is  poured 
over  a  strip  of  paper  and  is  then  immersed  in  a  jar  of  chlorine,  a 
violent  action  takes  place,  and  an  immense  cloud  of  finely  divided 
carbon  bursts  forth.  The  heat  of  the  action,  as  it  starts,  vaporizes 
the  turpentine,  and  the  hydrogen  in  the  latter  unites  with  the  chlorine 
forming  hydrogen  chloride,  while  the  carbon  is  set  free:  Ci0Hi6  + 
16C1  ->  16HC1  +  IOC. 

The  action  of  chlorine  on  potassium  iodide,  dry  or  in  solution,  is  of 
this  kind,  and  furnishes  the  commonest  test  (p.  66)  for  free  chlorine. 
The  chlorine  simply  takes  the  place  of  the  iodine  (KI  +  C1— »KC1  +  I), 
and  the  latter  is  liberated.  Iodine,  when  moist,  is  deep  brown  in 
color,  and  the  amount  liberated  by  a  large  quantity  of  chlorine  may 
easily  be  seen.  Yet,  it  is  an  advantage  so  to  arrange  a  test  that  it 
may  be  as  delicate  as  possible,  that  is,  may  give  a  plainly  visible 
result  with  the  minimum  of  material.  In  this  case,  much  starch 
emulsion  and  a  little  potassium  iodide  are  employed.  Strips  of  filter 
paper  dipped  in  this  mixture,  when  brought  into  a  gas  containing 
even  a  trace  of  free  chlorine,  show  a  deep-blue  color  (see  Iodine). 
Combined  chlorine,  as  in  a  chloride,  has  no  effect  upon  starch. 


CHLORINE  115 

3.  When  actions  like  that  on  turpentine  are  moderated  by  proper 
means,  the  decomposition  is  not  so  complete.  If  methane  (marsh 
gas,  CH4)  is  mixed  with  chlorine  and  exposed  to  sunlight,  a  slower 
action  occurs,  of  which  the  first  stage  consists  in  the  removal  of  one 
unit  weight  of  hydrogen  and  the  substitution  of  chlorine  for  it  accord- 
ing to  the  following  equation: 

CH4  +  2C1  -*  CH3C1  +  HC1. 

The  process  may  continue  further  by  the  substitution*  of  chlorine 
for  the  units  of  hydrogen  one  by  one  until  carbon  tetrachloride  (CC14) 
is  finally  formed. 

A  most  interesting  and  important  action  of  this  class  occurs  when 
chlorine  is  dissolved  in  water.  A  small  part  of  the  chlorine  interacts 
with  a  little  of  the  water,  the  element  being  substituted  for  one-half 
of  the  hydrogen: 

2C1  +  HOH  -»  HOC1  +  HC1.  (1) 

Only  traces  of  hydrogen  chloride  and  hypochlorous  acid  are  produced, 
however,  and  a  state  of  equilibrium  is  reached.  The  change  comes 
quickly  to  a  standstill,  because  the  products  interact  even  more 
vigorously  to  reproduce  chlorine  and  water: 

HC1  +  HOC1  <=>  2C1  +  H20,  (2) 

the  interaction  being  reversible  (p.  46).  This  interaction  of  chlorine 
and  water  (1),  slight  as  it  is,  is  of  importance  on  account  of  the 
instability  of  the  hypochlorous  acid  (q.v.)  which  it  produces.  When 
the  solution  is  exposed  to  sunlight,  the  hypochlorous  acid  decom- 
poses and  oxygen  gas  is  produced :  HC1O  — >  HC1  +  O.  Since  this 
removes  the  substance  on  whose  interaction  with  the  hydrogen 
chloride  in  (2),  the  reversal  of  (1)  depends,  the  latter  action  pro- 
ceeds under  continuous  illumination  gradually  to  completion. 

*  Substitution  resembles  displacement  (p.  68)  in  that  an  element  and  a 
compound  interact,  and  the  element  takes  the  place  of  one  unit  in  the  com- 
position of  the  latter.  In  the  above  action,  one  unit  of  chlorine  takes  the 
place  of  one  unit  of  hydrogen.  But  the  latter  is  not  liberated:  it  combines 
with  another  unit  of  chlorine.  Double  decomposition  the  action  is  not, 
because  elements  do  not  decompose.  The  name  used  is  intended  to  fix  the 
attention  on  the  compound  and  on  the  fact  that  one  unit  has  been  substituted 
for  another  in  it.  This  conception  is  a  favorite  one  in  the  chemistry  of  com- 
pounds of  carbon. 


116  COLLEGE    CHEMISTRY 

Hence  the  aqueous  solution  of  chlorine  must  be  kept  in  the  dark, 
since  otherwise  a  dilute  solution  of  hydrogen  chloride  alone 
remains. 

The  reader  should  note  here  the  displacement  of  the  equilibrium, 
a  chemical  one  in  this  case,  in  consequence  of  the  annulment  of  one 
of  the  opposing  tendencies  (p.  91).  Through  the  destruction  of 
the  hypochlorous  acid,  one  of  the  tendencies,  namely  that  repre- 
sented in  the  forward  direction  of  (2),  becomes  inoperative. 

4.  Chlorine  may  simply  add  itself  to  a  compound.  Thus,  one  of 
the  oxides  of  carbon,  carbon  monoxide  CO,  when  mixed  with  chlorine 
and  exposed  to  sunlight  gives  drops  of  a  volatile  liquid  (b.-p.  8.2°) 
known  as  phosgene  COC12. 

Chemical  Relations  of  the  Element.*  —  In  the  formation  of 
chlorides,  an  atomic  weight  of  chlorine  is  equivalent  to  one  atomic 
weight  of  hydrogen  or  of  sodium.  The  element  is,  therefore,  univa- 
lent  (p.  70).  It  never  shows  any  higher  valence  than  this,  save  in 
its  oxygen  compounds  (see  Chap.  xvi).  The  oxides  of  chlorine  inter- 
act with  water  to  give  acids,  and  the  element  is,  therefore,  to  be 
classed  as  a  non-metal  (p.  82).  It  belongs  to  that  group  of  the 
non-metals  called  the  halogens  (q.v.),  as  a  consideration  of  some 
others  of  its  relations  will  show  (see  Chap.  xiv). 

Uses  of  Chlorine.  —  Large  quantities  of  chlorine  are  manu- 
factured for  the  preparation  of  bleaching  materials  and  disinfecting 
agents.  In  disinfection,  the  minute  germs  of  disease  and  putrefac- 
tion are  acted  upon  either  by  the  chlorine  or  by  the  hypochlorous 
acid  formed  by  its  interaction  with  water,  and  instantly  their  life  is 
destroyed.  One  of  the  processes  for  the  extraction  of  gold  involves 
an  action  of  chlorine  gas  upon  the  material  after  it  has  received 
preliminary  treatment.  A  chloride  of  gold  is  formed,  which  can  be 
dissolved  out  of  the  matrix  by  means  of  water,  and  the  metallic  gold 
is  afterwards  precipitated  from  the  solution. 

*  In  accordance  with  the  distinction  that^must  be  drawn  (p.  20)  between 
the  element  as  a  variety  of  matter  in  combination,  and  the  elementary  sub- 
stance or  free  form  of  the  element,  and  to  avoid  a  common  source  of  con- 
fusion, we  shall  always  give  only  the  behavior  of  the  elementary  substance 
under  the  title  chemical  properties.  The  characteristics  which  distinguish 
the  compounds  of  the  element,  as  a  class,  from,  or  relate  them  as  a  class  to  the 
compounds  of  other  elements  will  then  appear  in  a  separate  section  under 
the  title  "  Chemical  relations  "  (see  pp.  157,  172). 


HYDROGEN   CHLORIDE  117 

HYDROGEN  CHLORIDE  HC1. 

Preparation  from  Sodium  Chloride.  —  As  we  have  seen,  the 
direct  union  of  hydrogen  and  chlorine  produces  a  gas,  hydrogen 
chloride  (HC1).  For  the  purpose  of  preparing  this  gas,  however, 
some  more  easily  managed  method  will  naturally  be  employed. 

In  commerce  large  quantities  of  hydrogen  chloride  are  obtained  as 
a  by-product  in  connection  with  the  manufacture  of  soda.  The  same 
materials  are  commonly  employed  in  the  laboratory.  Common  salt 
is  treated  with  concentrated  sulphuric  acid  at  a  gentle  heat.  Effer- 
vescence is  seen  and  hydrogen  chloride  is  given  off,  while  a  com- 
pound known  as  sodium  hydrogen  sulphate  (sodium  bisulphate  or 
acid  sulphate)  remains  behind.  The  action  is  represented  by  the 
equation : 

NaCl  +  H2SO4  ->  HC1  f  +  NaHSO4.  (1) 

The  Apparatus  used  in  the  preparation  of  chlorine  (Fig.  43)  may  be 
employed.  On  account  of  its  extreme  solubility,  the  gas  is  not 
washed  in  water,  however.  For  the  same  reason  it  must  be  collected 
by  downward  displacement  of  air,  or  over  mercury. 

The  above  statements  apply  to  the  action  when  no  high  tempera- 
ture is  employed.  If,  however,  a  larger  proportion  of  salt  is  used 
and  a  sufficiently  high  temperature  produced  by  artificial  heating, 
then  neutral  sodium  sulphate  is  formed  according  to  the  equation: 

2NaCl  +  H2S04  ->  2HC1  +  Na2S04.  (2) 

The  former  action  is  that  which  occurs  under  the  conditions  used  in 
the  laboratory.  The  latter  is  the  action  which,  with  furnace  heat, 
is  employed  in  commerce,  since  it  is  for  the  purpose  of  making  so- 
dium sulphate  (Na2SO4),  from  which  sodium  carbonate  is  afterwards 
to  be  prepared,  that  the  operation  is  undertaken.  The  hydrogen 
chloride  passes  through  a  tower,  down  which  water  trickles  over 
lumps  of  coke,  and  is  dissolved.  The  aqueous  solution  is  called 
hydrochloric  acid,  or,  in  commerce,  muriatic  acid. 

Interaction  of  Acids  and  Chlorides.  —  It  should  be  noted: 

1.  That   the  above-mentioned  actions  are  both  double  decom- 
positions. 

2.  That  when  any  acid  is  used  in  a  double  decomposition,  another 
acid  must  result,  by  transfer  of  the  radical  hydrogen  which  all  acids 


118  COLLEGE   CHEMISTRY 

contain.  For  example,  with  ammonium  nitrate  and  sulphuric  acid 
we  obtain  nitric  acid: 

(NH4)NO3  +  H2S04->HN03  +  (NH4)HS04. 

3.  That  hydrogen  chloride,  in  particular,  can  be  formed  by 
double  decomposition  of  any  acid  with  any  substance  containing  the 
chloride  radical  (Cl). 

To  make  hydrogen  chloride,  use  of  an  acid  that  can  be  obtained 
free  from  any  large  amount  of  water  gives  the  best  result  (see  below). 
Thus,  concentrated  phosphoric  acid  with  any  chloride  will  give  a 
change  parallel  to  those  described  above.  With  potassium  chloride 
this  action  would  be,  KC1  +  H3PO4  ->  HC1 1  *  +  KH2PO4  (primary 
potassium  phosphate). 

TJie  Kinetic  Hypothesis  Applied  to  the  Interaction  of  Sul- 
phuric Acid  and  Salt.  —  One  who  has  used  the  above-described 
method  for  making  hydrogen  chloride  without  reflection  would  not 
realize  the  complexity  of  the  machinery  by  which  the  result  is 
achieved.  The  means  are  apparently  very  simple.  Yet  the 
mechanical  features  of  this  experiment,  when  laid  bare,  are  extremely 
curious  and  interesting.  A  single  fact  will  show  the  possibilities 
which  are  concealed  in  it. 

.  If  we  take  a  saturated  solution  of  sodium  hydrogen  sulphate  in 
water  and  add  to  it  a  concentrated  solution  of  hydrogen  chloride  in 
water  (concentrated  hydrochloric  acid) ,  we  shall  perceive  at  once  the 
formation  of  a  copious  precipitate.  This  is  composed  entirely  of 
minute  cubes  of  sodium  chloride : 

NaHSO4  +  HC1  ->  H2SO4  +  Nad  J.*  (3) 

Now  this  action  is  nothing  less  than  the  precise  reverse  of  (1),  yet  it 
proceeds  with  equal  success.  In  fact,  this  chemical  interaction  is  not 
only  reversible  (pp.  46,  51,  115),  but  can  be  carried  to  completion  in 
either  direction.  It  is  only  in  presence  of  a  large  amount  of  water 
that  it  stops  midway  in  its  career  and  is  valueless  for  securing  a 
complete  transformation  in  either  direction: 

NaHSO4  +  HC1  <=±  H2SO4  +  NaCl. 
In  an  action  which  is  reversible,  if  the  products  remain  as  perfectly 

*  The  arrow  directed  downwards  indicates  elimination  of  a  substance  by 
precipitation,  that  directed  upwards,  escape  as  a  gas  or  solution  of  a  solid. 


HYDROGEN   CHLORIDE  119 

mixed  and  accessible  to  egch  other  as  were  the  initial  substances, 
their  interaction  will  continually  undo  a  part  of  the  work  of  the 
forward  direction  of  the  change.  Hence,  in  such  a  case  the  reaction 
must,  and  does,  come  to  a  standstill  while  as  yet  only  partly  accom- 
plished (cf.  p.  115);  but  this  was  not  the  case  with  actions  (1)  and 
(3).  Let  us  examine  the  means  by  which  the  premature  cessation 
of  each  was  avoided. 

In  equation  (1)  the  salt  dissolved  to  some  extent  in  the  sulphuric 
acid,  NaCl  (solid)  +±  NaCl  (dslvd),  and  so,  by  contact  of  the  two 
kinds  of  molecules,  the  products  were  formed.  On  the  other  hand, 
the  hydrogen  chloride,  being  insoluble  in  sulphuric  acid,  escaped  as 
fast  as  it  was  formed:  HC1  (dslvd)  <=±  HC1  (gas).  Hence,  in  that 
case,  almost  no  reverse  action  was  possible,  and  the  double  decom- 
position went  on  to  completion.  With  all  the  sodium  hydrogen 
sulphate  in  the  bottom  of  the  flask,  and  most  of  the  hydrogen 
chloride  in  the  space  above,  the  two  products  might  as  well  have 
been  in  separate  vessels  so  far  as  any  efficient  re-interaction  was 
concerned.  This  plan,  in  which  water  is  purposely  excluded,  forms 
therefore  the  method  of  making  hydrogen  chloride. 

In  equation  (3),  on  the  other  hand,  the  hydrogen  chloride  was 
taken  in  aqueous  solution,  and  was  kept  permanently  in  full  contact 
with  the  sodium  bisulphate.  It  had  therefore,  in  this  case,  every 
opportunity  to  interact  with  the  latter  and  no  chance  of  escape. 
Every  molecule  of  each  ingredient  could  reach  every  molecule  of 
the  other  with  equal  ease.  Furthermore,  the  sodium  chloride 
produced  as  a  result  of  their  activity  is  not  very  soluble  in  con- 
centrated hydrochloric  acid  (far  less  so  than  in  water),  and  so  it 
came  out  as  a  precipitate:  NaCl  (dslvd)  +±  NaCl  (solid).  But  this 
was  almost  the  same  as  if  it  had  gone  off  as  a  gas.  It  meant 
that  the  greater  part  of  the  salt  was  in  the  solid  form.  It  was 
in  a  state  of  fine  powder,  it  is  true.  But,  in  the  molecular  point  of 
view,  the  smallest  particle  of  a  powder  contains  millions  of  mole- 
cules, and  most  of  these  are  necessarily  buried  in  the  interior  of  a 
particle.  Thus  the  sodium  chloride  was  no  longer  able  to  interact 
effectively  molecule  to  molecule  with  the  other  product,  the  sulphuric 
acid.  Hence,  there  was  little  reverse  action  to  impede  the  progress 
of  the  primary  one.  Thus  (3)  is  nearly  as  perfect  a  way  of  liberating 
sulphuric  acid  as  (1)  is  of  liberating  hydrogen  chloride. 

This  discussion  is  given  to  illustrate  the  displacement  of  a  chemical 


120  COLLEGE    CHEMISTRY 

equilibrium,  and  to  explain  the  method  of  preparing  hydrogen  chlo- 
ride. It  also  throws  an  interesting  light  on  chemical  affinity,  however. 
Considering  action  (1),  by  itself,  we  might  reason  that  the  hydrogen 
chloride  was  formed  because  the  affinity  of  the  hydrogen  (H)  for 
chlorine  (Cl)  was  greater  than  for  the  sulphate  radical  (SO4).  But, 
if  we  did  so,  then  in  action  (3)  we  should  be  compelled  to  reason 
similarly  that  the  preponderance  of  affinity  was  just  the  opposite. 
In  point  of  fact,  no  conclusion  about  relative  affinity  can  be  drawn 
from  these  actions.  The  effects  of  affinity  are  here  entirely  sub- 
ordinated by  the  effects  of  a  purely  mechanical  arrangement,  depend- 
ing on  solubility.  When  the  activities  of  the  acids  are  properly 
compared,  hydrochloric  acid  is  found  to  be  considerably  more  active 
than  sulphuric  acid. 

Physical  Properties.  —  Hydrogen  chloride  is  a  colorless  gas, 
which  produces  a  suffocating  effect  when  breathed. 

Density  (H  =  1),  18.23  Crit.  temp.,  +  52° 

Weight  of  1  1.,  1.64  g.  Boiling-point  (liq.),  -  83.7° 

Sol'ty  in  Aq  (0°),  50,300  vols.  in  100         Melting-point  (solid),  -  110° 

The  gas  is  one-fourth  heavier  than  air.  On  account  of  its  great 
solubility  it  condenses  atmospheric  moisture  into  a  fog  of  drops  of 
hydrochloric  acid.  Both  in  the  gaseous  and  liquefied  states  it  is  a 
nonconductor  of  electricity.  Its  heat  of  solution  (p.  100)  is  17,400 
calories.  On  account  of  its  high  concentration,  the  saturated,  aque- 
ous solution  may  be  looked  upon  as  a  mixture  of  liquefied  hydrogen 
chloride  and  water. 

When  the  concentrated  aqueous  solution  is  heated,  it  is  the  gas  and 
not  the  water  which  is  driven  out,  for  the  most  part.  When  the  con- 
centration has  been  reduced  to  20.2  per  cent,  the  rest  of  the  mixture 
distils  unchanged  at  110°.  If  a  dilute  solution  is  used,  water  is  the 
chief  product  of  distillation  (about  100°),  but  gradually  the  boiling- 
point  rises,  and,  when  the  concentration  has  reached  20.2  per  cent 
once  more,  the  same  hydrochloric  acid  of  constant  boiling-point,  as  it 
is  called,  forms  the  residue. 

Chemical  Properties.  —  Hydrogen  chloride  is  extremely  stable, 
as  we  might  expect  from  the  vigor  with  which  the  elements  of  which 
it  is  composed  combine.  On  being  heated  to  a  temperature  of 
1800°  it  begins,  however,  to  dissociate  into  its  constituents. 


DROGEX    CHLORIDE 


.  IT  U 

121 


In  the  chemical  pomvi^^Bpr,  it  is  on  the  whole  rather  an  indif- 
ferent substance.     Hydrogen  chloride  (the  gas)  has  no  action  upon  I 
any  of  the  non-metals,  such  as  phosphorus,  carbon,  sulphur,  etc. 
Many  of  the  metals,  however,  particularly  the  more  active  ones,  such 
as  potassium,  sodium,  and  magnesium,  decompose  it.     Hydrogen  is 


* 
* 

I 


FIG.  44. 


FIG.  45. 


set  free,  and  the  chloride  of  the  metal  is  formed  (K  +  HC1— »  KC1  +  H) . 
Hydrogen  .chloride  unites  directly  with  ammonia  gas  to  form  solid 
ammonium  chloride  (HC1  +  NH3  — >  NH4C1).  The  liquefied  gas  .has 
the  same  properties. 

Composition. — The  proportion  of  hydrogen  to  chlorine  by  weight 
in  this  compound  is  1.008  :  35.45. 


122  COLLEGE   CHEMISTRY 

*The  proportion  by  volume  in  which  '%he  constituents  unite,  and 
*'  the  relation  of  this  to  the  volume  of  the  resulting  hydrogen  chloride, 
may  easily  be  shown  in  several  ways.  The  decomposition  of  the 
solution  of  hydrogen  chloride  in  water  by  means  of  the  electric  cur- 
rent proves  that  the  gases  are  liberated  in  equal  volumes.  Brownlee's 
apparatus  for  demonstrating  this  is  shown  in  Fig.  44.  The  central 
part  is  the  same  as  in  Fig.  21,  but,  when  the  three-way  stopcock  is 
closed,  the  gajpgo  to  right  and  left,  and  displace  the  liquid  in  two 
outside  tubes.  The  equal  rate  at  which  this  takes  place  on  both 
sides  proves  that  the  gases  are  generated  in  equal  volumes. 

In  order  to  ascertain  the  relation  between  the  volumes  of  the  con- 

fcstituents  and  that  of  the  product,  we  may  unite  the  gases  and  find 

out  whether  any  change  in  volume  occurs.     A  tube  with  thick  walls 

'^[Fig.  45)  is  filled  with  the  mixed  gases  obtained  by  electrolysis.     By 

dipping  one  end  of  the  .tube  under  mercury  and  opening  the  lower 

^topcock,  it  is  seen  that  no  gas  leaves  and  no  mercury  enters.     After 

Pthe  mixture  has  been  exploded,  by  the  light  from  burning  magnesium, 

the  same  test  is  repeated  with  the  same  result.     The  pressure  has 

therefore  remained  equal  to  that  of  the  atmosphere.     Hence  there 

has  been  no  change  in  volume  as  the  result  of  the  union.     It  appears 

therefore,  that: 

1  vol.  hydrogen  +  1  vol.  chlorine  — »  2  vols.  hydrogen  chloride, 
a  result  in  harmony  with  Gay-Lussac's  law  (p.  84). 

Chlorides.  —  The  chlorides  are  described  individually  under  the 
other  element  which  each  contains.  For  the  present  we  simply  add 
to  the  statement  on  p.  118  that  thej»ajority  of  the  chlorides  of  the 
metals  are  easily  soluble  in  water.  The  chief  exceptions  are  silver 
chloride  (AgCl),  mercurous  chloride  (calomel,  HgCl),  cuprous  chlo- 
ride (CuCl),  and  l$d  chloride  (PbCl2).  The  last  of  these  is  on  the 
border  line  as  regards  solubility.  An  appreciable  amount  dissolves 
in  cold  water,  and  a  considerable  amount  in  boiling  water  (see  Table 
of  solubilities,  inside  the  cover  at  the  front  of  this  book). 

The  various  modes  of  preparing  chlorides  are  enumerated  in  the 
next  section. 

Chemtoal  Properties  of  Hydrochloric  Acid.  — The  solution  of 
hydrogen  chloride  in  water  is  an  entirely  different  substance  in  its 
chemical  behavior  from  hydrogen  chloride.  It  is  strongly  acid,  turn- 


HYDROGEN    CHLORIDE  123 

ing  litmus  red.  The  gas  and  liquefied  gas  have  no  such  property. 
The  solution  conducts  electricity,  as  we  have  seen,  very  well,  and  is 
decomposed  in  the  process,  while  the  gas  and  the  liquefied  gas  are 
practically  nonconductors. 

Many  metals,  when  introduced  into  hydrochloric^icid,  displace  the 
hydrogen  (p.  65),  and  form  the  chloride  of  the  metal.  In  the  case 
of  zinc  the  action  was  represented  oy  the  equation: 

A  +  2HC1  ->  ZnCl2  +  2H.  * 
of  hydrogen  chloride  interacts  rapidly  with 
most  oxides  and  hydroxides  of  metals,  as,  for  example,  those  of  zinc: 

'ZnO  +  2HC1  ->  ZnOj  4>H20, 

2HC1  ->  Zn«2  -^  2H2O. 


* 


Here  no  free  hydrogen  is  obtained,  since  the  oxygen  in  the  oxide,  and 
the  hytlroxyl  in  the  hydroxide,  unite  witli  it  to  form  water.  In  each 
case,  however,  the  chloride  of  the  metal  is  obtained.  It  may  be 
noted,  in  passing,  that  all  acids  behave  in  a  similar  manner  towarcHI 
oxides  and  hydroxides,  giving  water  and  a  compound  corresponding 
to  the  chloride  (cf.  p.  1  1  1)  .  Dilute  sulphuric  acid,  for  example,  gives 
sulphates. 

In  the  two  preceding  paragraphs,  three  kinds  of  actions,  each  con- 
stituting a  different  mode  of  preparing  chlorides,  have  been  mentioned 
incidentally.  There  are  two  others  which  we  have  already  encoun- 
tered. The  simplest  is  the  direct  union  of  the  element  with  chlorine 
(Zn  -f  2C1  —  »  ZnCL,).  r=The  other  method  is  illustrated  in  the  case  of 
the  precipitation  of  silver  chloride  (p.  8).  Here  the  formation  of 
the  chloride  occurred  by  exchange  of  another  radical  for  chlorine 
(AgNO3  4-  NaCl->  AgCl  I  +  NaN03).  The  insoluble  chlorides  (p. 
122)  can  yRiade  conveniently  by  this  plan.  The  formation  of  the 
precipitates,  for  example  that  of  silver  chloride,  is.  used  as  a  test  for 
the  presence  of  a  soluble  chloride  in  the  solution. 

The  solution  of  hydrogen  Chloride  in  water  sold  in  commerce  is 
known  by  the  name  of  munKic  acid  (Lat.  muria,  brine).  It  is  a 
yellow  liquid  which  contains  a  number  of  impurities.  The  most  com- 
mon are,  ferric  chloride,  which  is  responsible  for  part  of  the  yellow 
tint,  some  yellow  organic  coloring  material,  arsenious  chloride,  and 
free  chlorine.  The^acid  frequently  gives  a  residue  when  evaporated, 
and  this  must  of  course  represent  some  impurity.  It  is  sometimes 


• 


124  'COLLEGE  CHEMISTRY 

adulterated  with  calcium  chloride,  since^he  price  obtained  depends 
*n  the  specific  gravity  of  the  solution,  and  this  may  be  raised  by 
dissolving  calcium  chloride  in  it. 

Classification  of  Chemical  Interactions  and  Exercises 
Thereon.  —  So  ftir  we  have  denned  ten  more  or  less  distinct  kinds  of 
chemical  change,  seven  differing»in  mechanism:  Combination  (p.  9), 
decomposition  (p.  10),  dissociation  (p.  81),  displacement  (p.  68),  substi- 
tution (p.  11. 5),  double  decomposition  (pp.  10,  68),  and  internal  rear- 
rangement (p.  10);  and  three  others:  oxidation  (pp.  52~75,  110,  112), 
reduction  (pp.  53*5),  aad  electrolysis  (pp.  13,  64).  Illustrations  of  all 

ft  one  of  these  will  be  found  in  the  present  chapter*     Some  actions 
ong  to  one  of  the  first  aeve$,  and  also  to  one  of  the  three  other 
sses.     The  ability  readify  to   classify  fach  phenomenon,  as  it 
mes  up,  requires  precisely  that  grasp  of  the  framework  of  the 
Bence  which  the  read^^kst  seek  speedily  to  attain.    For  e^mple, 
^let  him  classtfj^he  fS^dng  actions:     1-  action  of  potassium  on 
water;  2.   of  heat  on  potassium  chlorate;  3.   of  chlorine  on  metals; 
Pi.   of  chlorine  on  turpentine;  5.   of  chlorine  on  potassium  iodide; 
6.   of  chlorine  on  methane;  7.   of  carbon  monoxide  and  chlorine; 
8.   of  sunlight  on  hypochlorous  acid;  9.   of  sulphuric  acid  on  salt; 
10.    of   zinc  oxide  and  hydrochloric  acid;    11.   of  zinc  on  hydro- 
chloric acid. 

12.  In  the  interactions  of  potassium  permanganate  and  of  man- 
ganese dioxide,  respectively,  with  hydrochloric  acid,  what  fractions 
of  the  whole   chlorine  are  liberated?    What  are  the  commercial 
advantages  of  the  use  of  salt  and  sulphuric  acift  with  the  manganese 
dioxide? 

13.  In  view  of  the  explanations  given,  can  you  define  the  general 
nature  of  trfl  sufl^fees  (p.  Ill)  which  may  be  usetkto  oxidize 
hydrochloric  acid? 


I 


f 


* 


CHAPTER  XII 


MOLECULAR   WEIGHTS    AND    ATOMIC   WEIGHTS 


Aii 


AVOGADRO'S.  hyp^hesis  (p.  89)  has  proved  tobe  by  far  the 
most  suggestive  and  fruitful  of  all  the  conceptions  developed  from 
the  kinetic-molecular  hypothesis.  We  are  now  m  a  position  to 
discuss  several  of  its  most  important  applications.  To  speak  in 
terms  of  the  h|JK>thesis,  these  concern  more  particularly  the  measure- 
ment of  the  relative  weights  of  the  molecules  of  different  gaseous 
substances,  and  the  determination  of  the  most  convenient  magni- 
tudes for  the  chemical  unit  weights  (atomic  weights;  cf.  p.  36). 

Cleaning  of  Avogadro9 s  Hypothesis.  —  First,  we  must  und 
stand  clearly  whig!  is  implied  in  the  statement  that :  In  equal  volumes 
of  all  gases',  at  the  same  temperature  and  pressure,  there  are  equal 
numbers  of  molecules.  It  means  that,  for  instance,  at  100°  and  760 
mm.,  in  all  specimens  of  gases  the  average  spacing  of  the  molecules 
,is  identical.  This  condition  is  independent  of  the  nature  of  the  gas 
—  £or  example,  Whether  it  is  a  Simple  or  a  compound  substance,  like 
oxygen  and  carbon  dioxide  respectively,  or  a  mixture,  like  air.  It 
means  that  when,  at  some  fixed  temperature,  we  fill  thlBte  vessel 
with  a  number  of  different  gases  or  gaseous  mixtures  successively, 
the  number  of  molecules  th^ft  will  hold  at  a  pressu  A  say,  of  one 
atmosphere  will  always  be  tflPfame.  If  we  take  care  to  keep  tem- 
perature and  pres^Bre  the  same,  the  equality  in  the  number  of  mole- 
cules that  will  enter  the  jar  will  take  care  onBeTf  automatically. 
In  what  follows,  to  avoid  continual  repetition,  it  is  tp  be  assumed 
that  temperatures  and  pressures  are  equal  unless  the  contrary  is 
expressly  stated. 

MOLECULAR  WEIGHTS. 

Tlie  Relative  Weights  of  the  Molecules.  —  According  to  Avo- 
gadro's  hypothesis,  vessels  of  equal  size  filled  with  different  gases 
contain  equal  numbers  of  gaseous  molecules.  Now  equal  volumes  of 
different  gases  differ  very  markedly  in  weight,  or,  in  other  words, 

125 


126 


COLLEGE 


CHEMISTRY 


the  densities  of  various  known  gases  caver  a  wide  range  of  values. 
(Thus,  hydrogen  is  the  ligMcst  of  all,  chlorine  is  more  than  thirty- 
five  times,  mercuric  chloride  (corrosive  sublimate)  vapor  over  one 
hundred  and  thirty-four  times  as  heavy.  Since  these  different 
weights  of  equal  volumes  represent  the  weights  A?  equal  numbers  of 
molecules,  the  difference  must  be  due  to  the  differing  weightsrof  the 
molecules  themselves.  The  densities  of  gases,  therefore,  may  be 
taken  as  measures  of  the  relative  weights  of  their  individual  molecules. 
The  extreme  significance  of  this  inference  in  chemistry  will  appear 
as  we  elaborate  upon  it. 

The  various  scales  on  which  the  densities  of  gases  may  be  calcu- 
lated, such  as  the  weights  of  one  liter  of  each  gas,  or  the  weights  of 
volumes  equal  to  that  of  one  gram  of  air,  are  illustrated  in  the  first 
two  columns  of  the  following  table: 


Weigh!  of 
One  Liter,  0° 
and  760  mm. 

Density, 
Air=»l. 

Molecular 
Weight,  i 
Ox.  -3  2. 

Hydrogen  

0.090 

0.0696 

2.016 

Oxygen  

1  .  429 

1.105 

32.CO 

Chlorine 

3.166 

2  449 

70  90 

Hydrogen  chloride 

1.628 

1.259 

36.458 

Carbon  dioxide                              .    . 

1.965 

1.520 

44.  CO 

Water                 .    .           

JL8045 

18.016 

Mercury                       ^    .       .... 

^932 

6.908 

200.0* 

Mercuric  chloride       .        

•    12.097 

9.354 

270.90 

Air 

1  293 

1.00 

28  955 

* 

The  values  tor  water  (b.-p.  100°),  mercury  (b.-p.  357°),  and  mer- 
curic chloride  (b.-p.  300°)  are  measured  at  high  temperatures  and 
reduced  by  rule  (p.  62)  to  0°  and  760  mm.  All  the  numbers  in  the 
first  two  columns,  as  they  stand,  are  purely  physical  in  derivation. 
Those  in  the  second  column  are  obtained  from  those  in  the  first  by 
using  the  proportion: 

1.293  (wt.  1  1.  aii^  :  1.00  (air  =  1)  : :  wt.  of  1  1.  any  gas  :  x  (dens,  of 
that  gaf). 

The  last  column  will  be  explained  presently.  Since  the  numbers  in 
the  first  column  apply  to  equal  volumes  (1  1.),  and  those  in  the 
second  stand  in  constant  ratios  to  them,  the  weights  in  the  second 
column  represent  equal  volumes  also.  In  the  second,  the  volume 


p 

MOLECULAR    WEIGHTS  127 

is  T.sW  !•  The  values  in  eitfier  one  of  the  columns  represent  the 
relative  weights  of  the  molecules  of  the  various  substances  (see 
Exercise  1  in  this  chapter). 

f 

Molecular  Weights.  —  The  foregoing  section  shows  that,  pro- 
vided a  substance  is  a  gas  or  can  be  volatilized,  the  relative  weight 
of  its  molecules,  compared  with  those  of  other  volatile  substances, 
can  be  ascertained.  To  save  words,  the  relative  weight  of  the  mole- 
cules of  a  substance  is  called  the  molecular  weight  of  the  substance. 
Since  the  absolute  weights  of  molecules  cannot  be  determined,  the 
next  question  which  arises  is  as  to  the  choice  of  an  appropriate  unit, 
and  therefore  an  appropriate  scale  for  these  relative  weights,  or 
molecular  weights.  Now  the  numbers  already  given,  in  the  first 
two  columns  of*  the  table,  are  purely  physical  data  and,  as  they 
stand,  lack  direct  relation  to  chemical  facts.  A  set  of  chemical  num- 
bers is  required  for  chemical  purposes.  We,  therefore,  proceed  next 
to  Aow  how  the  required  relation  between  the  relative  weights  of 
molecules  and  chemical  facts  can  be  established. 

Chemistry  deals  with  chemical  combination,  and  most  substances 
are  compounds.  If  we  fix  our  attention,  then,  first,  on  compound 
substances  and  their  molecules,  we  perceive  at  once  that  the  mole- 
cules of  a  compound  substance  must  contain  two  or  more  elements 
in  definite  proportions  by  weight.  We  may  therefore  extend  the 
molecular  hypothesis  by  supposing  that  a  molecule  is  composed  of 
smaller  parts,  which  we  call  atoms.  The  atoms  will  be  elementary. 
Thus  the  molecule  of  the  compound  will  contain  one  or  more  atoms 
of  each  of  the  component  elements.  This  conception,  the  starting- 
point  of  the  atomic  hypothesis,  is  elaborated  in  the  next  chapter. 
Its  introduction  in  the  present  connection,  however,  at  once  suggests 
two  ideas.  In  the  first  place,  since  we  can  now  determine  the  relative 
weights  of  molecules,  we  should  also  somehow  be  able,  with  the  help 
of  the  combining  proportions,  to  determine  the  relative  weights  of 
the  atoms  of  the  elements.  If  these  relative  weights  of  atoms  are 
properly  determined,  then  the  weights  of  the  atoms,  when  added 
together,  should  give  the  weight  of  the  molecule.  Thus,  —  and  this 
is  the  second  idea,  and  the  one  of  most  immediate  use  to  us, — the  scale 
for  relative  weights  of  molecules  must  be  chosen  with  reference  to 
the  scale  for  relative  weights  of  the  atoms,  so  that  the  former  weights 
may  always  include  the  latter.  That  is  to  say,  the  molecular  weights 


COLLEGE.  CHEMISTRY 


must  be  based  upon  the  combining  ^A^its.  This  means  that  th< 
scale  must  be  such  that  no  molecule  of  a  compound  of  hydrogen  shal 
receive  a  value  so  small  that  the  proportion  of  hydrogen  in  it  is  le| 

-       ^"fcW"VJ>  TIL         .  -Ul-  '  Ji  1_  *      *  "  _  1^   M  f  " 


tfan  1.008.  Furthermore,  since  the  comflniing  weight  of  oxygen  n 
IB  standard  for  combining  proportions,  it  is  desirable  to  use  thi 
substance  as  basis  of  the  scale  of  molecular  weights.  Now,  as  w< 
shall  find  (see  p.  131),  it  turns  out  that,  to  avoid  obtaining  proper 
tions  of  hydrogen  less  than  unity,  we  are  compelled  to  take  the  scal< 
32  for  the  molecular  weight  of  oxygen.  Our  chemical  scale  of  densi 
ties  is  therefore  calculated  to  the  scale,  density  of  oxygen  =  32.  This 
then,  is  the  answer  to  the  problem  with  which  the  section  opened 
The  third  column  of  the  table  (p.  126)  shows  the  results  of  reealcu 
lating  the  densities  to  this  chemical  scale.  The  proportion  used  is 

Density  of  Ox.  :  Density  of  Substance  : :  32  :  x. 
Thus,  if  we  take  the  densities  from  the  first  column,  the  value  foi 
water  isfound  by  the  proportion,  1.429  :  0.8045  : :  32  :x  (=  It 
That  is,  we  multiply  the  weight  of  a  liter  of  the  gas  by  32/1.^ 
get  the  molecular  weight. 

Since  the  gram  is  the  unit  of  weight.  32  g.  of  oxygen,  or  18.016  g.  o: 
water  is  called  the  gram-molecular  weight  of  the  substance.  It  wil 
be  noted  that  32  g.  is  not  the  weight  of  a  molecule  of  oxygen.  It  L 
the  weight  of  a  very  large,  and  not  exactly  known  number  of  mole 
cules  of  oxygen.  But,  whatever  that  number  of  molecules  of  oxygei 
is,  18.016  g.  of  water  contains  the  same  number  of  molecules,  and  th< 
other  wef|Jits  in  the  same  column  are  weights  of  numbers  of  mole 
cules  equal  to  these.  The  term  gram-molecular  weight  being  some 
what  ponderous,  we  abbreviate  it  to  molar  weight,  and  still  further  t< 
mole.  Thus,  a  mole  of  chlorine  is  70.9  g.  of  the  element,  and  a  molt 
of  hydrogen  chlorkle  is  36.458  g.  of  the  compound. 

The  Gram- Molecular  (Molar)  Volume.  —  The  weights  in  th< 
last  column  of  the  table  (p.  126)  must  represent  equal  volumes  oj 
the  different  gases.  This  follows  from  the  fact  that  they  are  derivec 
from  the  values  in  the  first  column  by  multiplying  bj^aconstanl 
ratio  (32/1.429).  and  the  volume  in  the  first  column  ^J^f  1  Uter 
The  actual  dimension  of  this  volume  is  evidently  3^BK9  liters 
which  is  almost  exactly  22.39,  or  in  round  numbers  22.4  liters.  Thi* 
volume  at  OP  and  760  mm.  holds  32  g.  of  oxygen,  70.9  g.  of  chlorine 
44.00  g.  of  carbon  dioxide,  or,  in  fact,  the  molar  weight  of  any  gaseous 

* 


MOLECULAR    WEIGHTS  120 

substance.  It  is  called,  therefore,  the  gram-molecular  volume1  G.M.V.j 
or  the  molar  volume.  It  may  be  defined  as  that  volume  which  contains 
one  mole  v  gram-molecular  weight)  of  any  gas  at  0°  and  760  mm.  At 
other  temperatures  and  pressures  the  G.M.V.  has  correspondingly 
different  val 

\j  The  G.M.V.  gives  us  a  concrete  conception  of  a  molar  weight. 
This  volume  is  represented  by  a  cul^e  <Fig.  46)  28.19  cm.  (or  about 
11.1   inches)    high.     Like    any   other 
volume,  it  holds  the  identical  numbers 
of  molecules  of  different  g 
capacity   at  0°  and  760  mm.  is  the 


number  of  molecules  in  32  g.  of  oxygen. 
Hence,  in  terms  of  the  hypothesis,  the 
weight  of  any  gas  which  fills  it  bears  to     — 
the  same  ratio  as  the  weight  of  a          • 


molecule  of  that  gas  to  the  weight  of  a 
molecule  of  oxygen.   *  We  may,  there- 
fore, state  the  method  of  finding  the  molar    gram-molecular)  weight  of 
a  substance  thus:  Weigh  a  known  volume  of  the  substance,  at  any  tem- 
perature and  pressure  at  which  it  is  gaseous,  reduce  this  volume  by  rule 
to  0°  and  760  mm.,  and  calculate  by  proportion  the  weight  of  22.4  liters 

Exercise  5). 

That  quantity  of  each  substance  which  at  0°  and  760  mm  would  fill 
the  G.M.V.  cube  is  the  unit  quantity  of  the  substance  for  all  theoretical 
purposes  in  chemistry.  It  represents  the  relative  weight  of  the  mole- 
cules of  the  substance.  We  shall  employ  it  at  once  for  the  purpose 
of  determining  the  relative  weights  of  atoms,  or  atomic  weights. 

*  A  common  question  is:  Do  Dot  molecules  of  different  substances  differ 
in  size,  and  will  not  the  numbers  required  to  fill  the  GJi.Y.  therefore  be  differ- 
ent? The  answer  is  that  the  molecules  are  all  so  small  compared  with  the 
spaces  between  them  (at  760  mm.)  that  the  distances  from  surface  to  surface 
are  practically  the  same  as  from  center  to  center.  A  G.M.V.  of  oxygen, 
when  liquefied,  gives  less  than  32  c.c.  of  liquid  oxygen,  or  leas  than  1/700  of 
volume  as  gas.  And  there  are  still  spaces  between  the  molecules  of  the 
liquid.  It  is  only  when  gases  are  so  severely  compressed  that  the  nearness  of 
the  molecules  to  one  another  approaches  that  found  in  the  liquid  condition 
the  elfe&s  of  the  bulk  of  the  molecules  becomes  conspicuous,  and  a 
difference  in  the  behavior  of  different  gases  is  noticeable.  But  in  the  work 
discussed  in  this  chapter,  pressures  over  one  atmosphere  are  not 


130  COLLEGE   CHEMISTRY 


ATOMIC  WEIGHTS 

Determination  of  the  Atomic  Weight  of  EacJi  Element.  — . 

If  the  paragraphs  dealing  with  combining  weights  are  now  re-read 
(pp.  31-38),  it  will  be  found  that  the  foundations  for  a  system  of 
weights  was  worked  out,  but  that  no  basis  for  definitely  fixing  the 
individual  values  was  discovered.  At  the  time,  the  only  information 
we  had  was  obtained  by  analyzing  compounds  and  reasoning  about 
the  results,  and  evidently  something  more  was  needed  for  the  abso- 
lute determination  of  the  values.  Thus  on  p.  35,  the  equally 
good  standing  of  two  different  possible  values  for  the  combining 
weight  of  copper,  namely  63.6  and  31.8,  was  brought  out.  The 
same  dilemma  would  have  occurred  with  every  element  which  com- 
bines in  two  or  more  different  proportions  with  any  other,  and  there- 
fore affects  nearly  the  whole  list.  In  the  history  of  chemistry,  the 
same  uncertainty  caused  much  controversy  among  chemists  and,  for 
years,  endless  confusion.  It  was  only  when  the  comparison  of  the 
combining  weights  with  the  relative  weights  of  molecules  was  at 
last  rigorously  applied,  in  the  fashion  now  to  be  shown,  that  perfect 
order  in  chemical  weights  and  formulae  was  finally  achieved. 

To  determine  the  atomic  weights,  the  plan  of  procedure  is  per- 
fectly simple.  In  the  preceding  section  we  settled  upon  the  chemical 
unit  quantity  of  each  substance.  This  is  the  quantity  which,  in  the 
gaseous  condition,  would  fill  the  G.M.V.  (22.4  liters)  at  0°  and  760 
mm.  Now,  we  seek  the  chemical  unit  quantities  of  the  elements 
combined  in  each  substance.  Evidently  the  logical  and  consistent 
plan  must  be  to  take  the  amount  of  each  substance  which  fills  the 
G.M.V.  and  find  out  how  much  of  each  element  present  is  contained 
in  this  unit  amount  of  the  substance.  In  other  words,  to  put  the 
matter  concretely,  we  imagine  ourselves  filling  the  cube  (Fig.  46) 
with  one  compound  after  another,  and  in  each  case  determining  by 
analysis  the  weight  of  each  constituent  element  present  in  a  cube- 
full  of  the  substance.  To  carry  out  this  plan,  two  experimental 
operations  are  necessary  with  each  substance : 

First  we  determine  the  density,  and  this  gives  us  the  gram-molec- 
ular weight,  i.e.  the  amount  filling  the  cube.  This  shows  the  rela- 
tive weight  of  a  molecule  of  the  substance,  as  compared  with  that 
of  one  molecule  of  oxygen. 

Then  we  analyze  the  substance,  and  this  gives  us  the  quantity 


ATOMIC    WEIGHTS  131 

of  each  constituent  in  the  total  gram-molecular  weight,  i.e.  in 
the  material  filling  the  cube.  This,  in  turn,  shows  the  weight  of  the 
quantity  of  each  element  present  in  each  molecule,  relative  to  the 
weight  of  a  molecule  of  oxygen. 

For  example,  the  cube  holds  36.458  g.  of  hydrogen  chloride,  and 
this  amount,  when  decomposed,  yields  1.008  g.*  of  hydrogen  and 
35.45  g.  of  chlorine. 

Finally,  to  determine  the  best  combining  weight  for  a  given  ele- 
ment, we  repeat  the  two  foregoing  operations  with  as  many  different 
compounds  of  the  element  as  possible,  and  then  we  examine  the 
various  quantities  of  the  element  found  in  the  G.M.V.  of  the  various 
compounds.  From  inspection  of  these  quantities  we  quickly  select 
the  value  of  which  all  are  multiples,  by  unity  or  some  integral  number. 
This  value  for  the  combining  weight  is  the  one  accepted.  In  terms 
of  the  hypothesis,  this  is  the  weight  of  one  atom  of  the  element,  com- 
pared with  the  weight  of  a  molecule  of  oxygen,  and  molecules  con- 
taining more  than  this  proportion  contain  two,  three,  or  more  atoms 
of  the  element. 

For  example,  if  we  are  seeking  the  atomic  weight  of  chlorine,  we 
set  down  the  result  for  hydrogen  chloride  just  given.  Then  we  take 
another  compound  of  chlorine,  say  phosphorus  oxy chloride.  We 
determine  the  weight  of  a  measured  volume  of  its  vapor,  at  a  prop- 
erly chosen  temperature  and  pressure,  and  the  result  gives  us,  by 
calculation,  the  molecular  weight,  viz.  153.35.  That  is,  153.35  g. 
of  the  substance  would  fill  the  cube,  if  it  could  be  kept  as  vapor  at 
0°  and  760  mm.  And  this  amount  of  the  substance  contains  31  g. 
of  the  element  phosphorus,  16  g.  of  the  element  oxygen,  and  106.35 
g.  of  the  element  chlorine.  We  then  continue  the  processes  described, 
using  all  the  volatile  compounds  of  chlorine.  The  involatile  com- 
pounds (like  common  salt)  must  be  set  aside,  for  they  cannot  be 
vaporized,  and  therefore  their  molecular  weights  cannot  be  deter- 
mined. When  we  have  studied  as  many  compounds  as  possible  in 
this  way,  we  find  that  there  are  different  quantities  of  chlorine  in 
our  list,  but  they  are  all  integral  multiples  of  35.45  g.  In  phosphorus 

*  It  will  be  observed  that  if  the  unit  for  molecular  weights  had  been  less 
than  the  number  of  molecules  in  32  g.  of  oxygen,  then  an  equal  number  of 
molecules  of  hydrogen  chloride  would  have  contained  less  than  1.008  g.  of 
hydrogen,  and  the  atomic  weight  of  this  element  would  then  have  been  less 
than  unity. 


132 


COLLEGE   CHEMISTRY 


oxychloride,  for  example,  the  quantity  was  106.35,  or  3  X  35.45. 
Hence  35.45  g.  can  be  taken  as  the  unit  quantity,  the  atomic  weight 
of  the  element  chlorine.  In  terms  of  the  hypothesis,  this  is  the 
relative  weight  of  an  atom  of  chlorine,  as  compared  with  that  of  a 
molecule  of  oxygen,  when  the  value  32  is  assigned  to  the  latter.* 

In  the  following  table  a  few  sample  results  of  the  process  just  out- 
lined are  given.  The  first  column  contains  the  molar  weight,  i.e. 
the  weight  of  the  substance  which  occupies  the  G.M.V.  cube.  In 
the  other  columns  are  entered  the  weights  of  the  various  elements 
which  together  make  up  the  total  molar  weight.  To  simplify  the 
numbers,  the  value  1  is  used  for  hydrogen,  instead  of  1.008. 


Substance. 


Hydrogen  chloride     .    . 
Chlorine  dioxide     .    .    . 
Phosphorus  trichloride 
Phosphorus  oxychloride 
Phosphoric  anhydride  . 

Phosphine 

Water 

Methane 

Acetylene 

Ethylene 

Formaldehyde     .... 

Acetic  acid 

Mercurous  chloride    .    . 
Mercuric  chloride   . 


Weights  of  Constituents  in 
Mohir  Weight. 

Molar 
Weight 

c 

02 

Chlo- 

c' 

O 

£ 

Molecular 

rine. 

§i 

0, 

o 

g 

Formula. 

w 

O 

1 

1 

fc 

36.45 

1 

35.45 

HC1 

67.45  ... 

35.45    32 

C1O2 

137.35 

106.35 

31 

... 

PC13 

153.35 

106.35 

16 

31 

POC13 

284 

160 

124 

P4O10 

34 

3 

31 

PH3 

18 

2 

16 

H9O 

16 

4 

. 

12 

CH4 

26 

2 

24 

.  .  . 

C2H, 

28 

4 

24 

C;H; 

30 

2 

16 

12 

CH.O 

60 

4 

32 

24 

C9H4O2 

235.45 

35.45 

200 

HgCl 

270.9    .... 

70.9 

200 

HgCl2 

i 

*  It  should  be  noted  that  there  is  another  unit  quantity  of  chlorine, 
namely  the  molecular  weight,  or  weight  of  the  G.M.V.  of  the  substance.  This 
is  the  unit  quantity  of  free  chlorine.  But  we  are  dealing  now  with  compounds, 
and  proportions  in  combination,  so  that  free,  uncombined  chlorine,  and  other 
elements  in  free  condition  do  not  interest  us  at  present,  and  will  be  taken  up 
later. 


ATOMIC    WEIGHTS  133 

To  contain  similar  data  for  all  the  volatile  compounds  of  every 
known  element,  a  huge  table,  of  which  this  might  be  a  small  corner, 
would  be  required.  With  such  a  table  at  hand  the  atomic  weight 
of  each  element  could  promptly  be  picked  out.  Thus,  in  the  carbon 
column  it  would  be  found  that  all  the  weights  of  carbon  were  either 
12  or  integral  multiples  of  12,  and  this  is  therefore  the  atomic  weight 
of  carbon.  Similarly  the  atomic  weight  of  oxygen  is  16,*  of  phos- 
phorus 31,  of  mercury  200  (see  Exercise  4). 

When  the  atomic  weights  have  finally  been  selected,  we  can  go 
through  the  table  and  change  all  the  numbers  into  multiples  of  the 
chosen  atomic  weights.  Thus,  for  70.9  we  write  2  X  35.45,  and  for 
106.35  we  write  3  X  35.45,  and  so  forth.  The  reader  can  prepare 
such  a  modification  of  the  table.  With  this  new  form  of  the  table 
before  us,  we  can,  finally,  replace  the  atomic  weights  by  the  symbols 
which  stand  for  them,  writing,  for  35.45,  Cl,  for  2  X  35.45,  C12,  and 
so  forth.  The  results  of  doing  this  in  each  line,  i.e.  for  each  sub- 
stance, are  collected  at  the  ends  of  the  lines  in  the  last  column  of 
the  table.  The  reader  should  himself  repeat  the  substitutions  of  the 
symbols,  and  so  verify  the  formulae  given.  These  formulae,  since 
they  are  based  on  the  molecular  weights,  in  such  a  way  that  when 
the  numerical  values  are  substituted  for  the  symbols  the  total  re- 
stores to  us  the  molecular  weight,  are  called  molecular  formulae. 

It  will  now  be  seen  why  the  equivalents  (pp.  36-37)  were  multiplied 
by  various  integers  in  making  the  chemical  units.  The  equivalent 
of  carbon  was  3.  That  is  to  say,  carbon  and  oxygen  combine  in  the 
ratio  3:8  (in  carbon  dioxide),  and  carbon  and  hydrogen  in  the 
ratio  3  :  1.008  (in  methane).  But  there  is  no  compound  of  carbon 
whose  molecular  weight  contains  less  than  12  parts  of  the  element. 
It  would  thus  lead  to  needless  complication  to  take  3  as  the  unit 
amount  of  carbon,  for  every  molecule  would  then  contain  four  units, 
or  some  multiple  of  four,  and  every  formula  C4  or  some  multiple  of 
C4.  We  choose  the  largest  units  of  combining  weight  that  we  can, 
in  order  that  the  coefficients  may  be  the  smallest  possible,  and  the 
resulting  formulae  the  simplest  possible.  Naturally  the  actual  ratios 
remain  the  same.  Thus,  for  carbon  dioxide  the  ratio  3  :  8  is  replaced 
by  12  :  32,  or  12  :  2  x  16,  or  C  :  2O,  which  has  the  same  value. 

*  The  difference  between  the  unit  quantity  of  oxygen  in  compounds 
(namely,  16)  and  the  unit  quantity  of  free  oxygen  (32)  will  be  discussed 
presently. 


134  COLLEGE    CHEMISTRY 

As  a  definition,  the  atomic  weight  of  an  element  may  be  stated  to 
be:  The  smallest  of  the  weights  of  the  element  found  in  the  molecular 
weights  of  all  its  volatile  compounds,  so  far  as  these  have  been  exam- 
ined. Since  the  atomic  weight  is  always  a  multiple  of  the  equiva- 
lent weight  (by  unity,  or  some  other  integer),  it  might  also  be 
defined  as:  The  largest  even  multiple  of  the  equivalent  which  can 
be  contained  in  the  molecular  weights  of  all  the  volatile  compounds 
of  the  element.  The  complete  list  of  accepted  atomic  weights  is 
printed  on  the  inside  of  the  cover  at  the  back  of  this  book. 

- 

Advantages  of  Atomic  Weights  over  Equivalents.  —  Since  the 
method  of  determining  atomic  weights  depends  on  rather  complex 
reasoning,  and  involves  much  experimental  work,  the  question  may 
be  asked  whether,  when  found,  they  are  worth  all  the  trouble.  It 
is  manifest  that  equivalents  are  much  simpler  in  nature,  and  much 
more  easily  ascertained  than  atomic  weights.  It  will  be  expected, 
therefore,  that  we  shall  be  able  to  show  that  the  units  Na  =  23, 
Cu  =  63.6,  Al  =  27.1,  C  =  12,  etc...  give  a  better  view  of  the  rela- 
tions of  the  elements  than  do  the  equivalents  23,  31.8,  9.03,  and  3, 
respectively.  Now,  the  atomic  weights  are  as  good  as  equivalent 
weights,  for  the  purpose  of  acting  as  units,  in  terms  of  which  to 
express  combining  proportions,  and  possess,  besides,  several  (at 
least  five)  important  properties  or  uses  which  equivalent  weights 
entirely  lack. 

Of  these  valuable  properties,  the  first  two  have  been  mentioned 
already: 

1.  Being  very  often  themselves  larger  numbers  than  the  equiva- 
lents, atomic  weights  are  often  multiplied  by  smaller  coefficients 
(p.  133).     This  simplifies  our  equations. 

2.  The  atomic  weight  of  an  element  can  have  but  one  value,  and 
is  definitely  determinate.     Most  equivalents  have  more  than  one 
value  (p.  33). 

3.  The  atomic  weight  of  an  element  has  a  valence  (p.  68).  while 
equivalents  are  equi-valent.     While  valence  is  a  helpful  conception 
in  all  branches  of  chemistry,  organic  chemistry  is  especially  indebted 
to  the  conception  of  the  quadrivalence  of  carbon  for  much  of  its 
development  and  most  of  its  organization.     The  full  illustration 
of  this  point  is  beyond  the  limits  of  the  present  book. 

4.  The  periodic  system  (q.v.),  the  basis  of  a  plan  for  classifying 


ATOMIC   WEIGHTS 


135 


the  properties  of  all  chemical  substances,  is  founded  upon  the  atomic 
weights. 

5.  D along  and  Petit 's  law  is  based  upon  atomic  weights.  This  law 
furnishes  also  an  alternative  means  of  determining  atomic  weights 
that  has  frequently  rendered  valuable  service,  and  on  this  account 
forms  the  subject  of  the  next  section. 

Dulong  and  Petit  9s  Law,  an  alternative  Means  of  Deter- 
mining Atomic  Weights.  —  It  was  first  pointed  out  (1818)  by 
Dulong  and  Petit,  of  the  Ecole  Polytechnique  hi  Paris,  that  when 
the  atomic  weights  of  the  elements  were  multiplied  by  the  specific  heats 
of  the  simple  substances  in  the  solid  condition,  the  products  were 
approximately  the  same  in  all  cases.  In  other  words,  the  specific 
heats  are  inversely  proportional  to  the  magnitudes  of  the  atomic 
weights.  The  table,  in  which  round  numbers  have  been  used  for 
the  atomic  weights,  shows  that  the  product  lies  usually  between  6 
and  7,  averaging  about  6.4: 


Element. 

Atomic 
Wt. 

Sp.  Ht. 

Pro- 
duct. 

Element. 

Atomic 
Wt. 

Sp.  Ht. 

Pro- 
duct. 

Lithium 

7 

94 

6  6 

Iron 

56 

112 

6  3 

Sodium 

23 

29 

6  7 

Zinc 

65  4 

093 

6  1 

Magnesium    .    . 

24.4 

.245 

6.0 

Bromine  (Solid) 

80 

.084 

6.7 

Silicon    .... 

28.4 

.16 

4.5 

Gold  

197 

.032 

6.3 

Phosphorus 
(Yellow) 

31 

.19 

5.9 

Mercury 
(Solid) 

200 

.0335 

6.7 

Calcium     .    .    . 

40 

.170 

6.8 

Uranium  .    .    . 

238.5 

.0276 

6.6 

Another  way  of  expressing  this  law  will  give  it  greater  chemical 
significance.  The  specific  heats  are  the  amounts  of  heat  required  to 
raise  equal  weights  of  the  various  elements  through  one  degree. 
Now  these  equal  weights  contain  fewer  chemical  units  in  proportion 
as  the  chemical  unit  weight  is  greater.  Henc*  this  law  may  be  put 
in  the  form:  Equal  amounts  of  heat  will  raise  atomic  weights  of  all 
elements  through  equal  intervals  of  temperature. 

This  being  true,  the  equivalents,  if  used  instead  of  the  atomic 
weights,  must  give  widely  varying  products.  The  quantities  of  heat 
required  to  raise  equivalent  weights  through  one  degree  are  either 
equal  to,  or  are  fractions  of,  those  required  for  the  atomic  weights, 


136  COLLEGE    CHEMISTRY 

according  to  the  valence  of  the  element.     Hence  the  law  applies 
only  to  atomic  weights,  and  not  to  equivalents. 

It  will  be  seen  at  once  that  although  the  law  of  Dulong  and  Petit 
is  purely  empirical,  it  may  nevertheless  be  used  for  fixing  the  atomic 
weight  of  an  element  of  which  no  volatile  compounds  are  known.  We 
can  always  measure  the  equivalent  with  considerable  exactness,  and, 
when  this  has  been  multiplied  by  the  specific  heat  of  the  free  sub- 
stance, we  can  see  at  a  glance  what  integral  factor  will  raise  the 
product  to  the  neighborhood  of  6.4.  For  example,  analysis  shows 
us  that  in  calcium  chloride  the  proportion  of  chlorine  to  calcium, 
using  the  known  atomic  weight  of  chlorine  as  one  term  of  the  pro- 
portion, is  35.5  :  20.  If  calcium  is  univalent,  20  is  its  atomic  weight. 
If  it  is  bivalent,  two  units  of  chlorine  are  combined  with  40  parts  of 
calcium,  and  40  is  its  atomic  weight.  If  it  is  trivalent,  three  units  of 
chlorine  are  united  with  60  parts  of  calcium,  etc.  All  we  learn  in 
reference  to  the  atomic  weight  of  calcium  from  this  analysis  is  that 
its  value  is  20  or  some  integral  multiple  of  20.  Nor  can  we  fix  the 
upper  limit,  for  we  are  unable  to  obtain  the  weight  of  a  known 
volume  of  calcium  chloride  vapor  and  so  determine  the  molecular 
weight.  But  the  specific  heat  of  solid  calcium  being  0.170,  we  mul- 
tiply this  number  by  20,  and  get  the  product  3.4.  This  is  only  half 
large  enough,  so  we  assume  that  40  is  a  more  probable  value  for  the 
atomic  weight  of  calcium.  The  product  is  then  6.8,  which  agrees 
fairly  well  with  the  average  for  other  elements.  We  decide,  there- 
fore, that  the  symbol  Ca  shall  represent  forty  parts  by  weight.  The 
formula  of  calcium  chloride  is  therefore  CaCl2,  and  calcium  is  bivalent. 

MOLECULAR  FORMULA 

Molecular  Formulce  of  Compounds.  —  If  the  molar  formulae 
in  the  table  (p.  132)  be  examined  it  will  be  observed  that  several  are 
not  in  their  simplest  terms.  Thus,  the  formula  of  acetylene  is  C2H2. 
The  formula  CH  would  represent  the  composition  of  the  substance 
equally  well,  for  12  :  1  is  the  same  as  24  :  2.  But  the  formula  CH 
gives  a  total  of  only  13,  while  C2H2  shows  the  total  weight  of  the 
molecule  to  be  26  and  records  for  us  therefore  the  weight  of  the  G.M.V., 
as  well  as  the  composition  of  the  substance.  And  we  shall  find  this 
additional  property,  peculiar  to  the  molecular  formula,  to  be  a  feature 
of  the  greatest  practical  value.  Some  of  the  practical  uses  of  this 
improvement  in  our  formula  will  be  illustrated  in  this  chapter,  and 


MOLECULAR   EQUATIONS  137 

there  is  an  example  of  one  of  them  in  the  table  itself.  Thus,  the 
molecular  formula  of  acetic  acid  is  C2H4O2,  and  not  the  simpler, 
identical  proportion  CH2O.  The  latter  is  the  molecular  formula  of 
a  totally  different  substance,  formaldehyde,  now  much  used  as  a 
disinfectant.  The  vapor  of  this  substance  has  only  half  the  density 
of  acetic  acid  vapor,  and  this  fact,  recorded  in  the  formula,  helps  to 
remind  us  that  the  substances  are  different.  Still  another  substance 
of  the  same  composition  is  fruit  sugar  (dextrose),  CgH^Oe  (see  Exer- 
cise 12).  In  addition  to  this  and  other  practical  advantages,  molec- 
ular formulae  satisfy  also  the  claim  of  logical  consistence.  If  the 
symbols  represent  the  atomic  weights,  the  formulae  should  be  con- 
structed so  as  to  represent  the  molecular  weights. 

Molecular  formulas  like  C2H2  and  C2H4O2  are  easily  interpreted  in 
terms  of  the  atomic  hypothesis.  C  represents  one  atom  of  carbon 
and  H,  one  atom  of  hydrogen.  But  there  is  no  reason  why  a  mole- 
cule of  acetylene  should  not  contain  two  atoms  of  each  kind.  Simi- 
larly, the  molecule  of  formaldehyde  contains  four  atoms  (CH2O),  and 
one  of  acetic  acid  eight  atoms  (C2H402),  and  one  of  dextrose  twenty- 
four  atoms  (C6H12O6),  although  the  relative  numbers  of  each  kind 
are  the  same.  Indeed  this  hypothesis  helps  to  clear  the  matter  up, 
for  chemists  go  so  far  as  to  account  for  the  chemical  behavior  of  the 
substances  by  an  imagined  geometrical  arrangement  of  the  atoms  in 
their  molecules,  and  these  three  kinds  of  molecules  are  supposed  to 
differ  in  structure  as  well  as  in  the  number  of  atoms  they  contain. 

The  Molecular  Weights  and  Formulas  of  Elementary  Sub- 
stances.— The  following  table  gives  the  densities  of  some  elementary 
substances,  including  those  of  which  the  substances  last  discussed  are 
compounds.  The  first  column  shows  the  atomic  weight,  which  in 
each  case  is  the  minimum  weight  of  the  element  found  in  a  G.M.V. 
of  any  compound.  For  example,  16  g.  of  oxygen  and  35.45  g.  of 
chlorine  are  the  weights  in  the  amounts  of  water  vapor  and  hydrogen 
chloride,  respectively,  which  fill  the  cube  (22.4  liters).  The  symbol, 
in  the  next  column,  stands  for  this  quantity  and  occurs  in  many 
formulae,  such  as  H20  and  HC1.  It  represents  the  combining  unit 
or  atom.  In  the  third  column  is  given  the  density  of  the  free,  ele- 
mentary substance.  This  weight  of  the  simple  substance  fills  the 
G.M.V.  and  is  the  molecular  weight.  It  shows  the  weight  of  the 
molecule  relative  to  the  weights  of  the  other  molecules  in  the  same 


138 


COLLEGE    CHEMISTBY 


column,  and  to  the  weights  of  the  atoms  in  the  first  column.  In  the 
last  two  columns  are  given  the  densities  resolved  into  multiples  of 
the  atomic  weights  and  the  corresponding  formulae. 


Atomic 
Weight. 

Sym- 
bol. 

Density 
O  =»  32. 

Density  Fac- 
torized. 

Formula 
ol  Free 
Element. 

Oxvsren 

16.00 

o 

32.00 

2X16.00 

O, 

Hydrogen  

1.008 

H 

2.016 

2X1.008 

H2 

Chlorine     

35.45 

Cl 

70.90 

2X35.45 

C12 

Phosphorus       

31.0 

P 

124.0 

4X31.0 

P42 

Mercury     
Ozone     

200.0 
16.00 

%s 

200.0 
48.00 

1X200.0 
3X16.00 

Hg 
O, 

Cadmium 

112  4 

Cd 

112  4 

1X112  4 

Cd 

Pot&ssium 

39.15 

K 

39  15 

1X39  15 

K 

Sodium 

23.05 

Na 

23.05 

1X23  05 

Na 

Zinc 

65  4 

Zn 

65.4 

1X65  4 

Zn 

The  reader  cannot  fail  to  note  a  striking  peculiarity.  In  the  case 
of  chlorine  the  molecular  weight  is  70.9,  while  the  atomic  weight  is 
35.45.  With  hydrogen  and  oxygen,  also,  the  molecular  weight  con- 
tains two  atomic  weights.  Yet  this  is  not  a  general  rule,  for  with 
mercury  and  several  other  elements  the  molecular  and  atomic 
weights  are  alike,  while  with  phosphorus  the  molecular  is  four  times 
the  atomic  weight.  Evidently  there  is  no  rule,  and  each  element 
has  to  be  subjected  to  separate  experimental  study.  The  result  is 
that  for  free,  elementary  chlorine  we  use  the  molecular  formula  C12,  for 
free  hydrogen  H2,  for  elementary,  uncombined  oxygen  the  formula  02. 
For  a  substance  like  phosphorus,  which  is  not  a  gas  and  is  not  often 
used  as  a  vapor,  the  formula  P  is  commonly  employed  by  chemists, 
to  avoid  the  larger  coefficients  which  P4  introduces  into  equations, 
although  theoretically  the  latter  formula  would  be  the  strictly  correct 
one. 

The  case  of  oxygen  demonstrates  clearly  the  necessity  of  using 
molecular  formulae,  even  for  simple  substances.  The  table  shows 
two  substances  containing  nothing  but  oxygen.  Ozone  (q.v.)  has 
a  molecular  weight  48,  being  a  gas  exactly  one-half  heavier  than 
ordinary  oxygen.  Its  formula,  therefore,  is  08,  while  that  of  oxygen 
is  02.  Oxygen  and  ozone  are  entirely  different  chemical  individuals. 
The  latter  has,  for  example,  a  strong  odor  and  is  much  more  active. 


MOLECULAR  EQUATIONS  139 

Thus  polished  silver  remains  bright  indefinitely  in  pure  oxygen, 
!>iif  oxidizes  quickly  when  phc^d  in  ozone. 

To  avoid  a  common  error,  the  reader  should  note  that  to  learn  the 
atomic  weight  of  an  element,  we  do  not  measure  the  molecular  weight 
of  the  simple  substance.  The  molecular  weight  of  the  elementary 
substance  may  be  a  multiple  of  the  atomic  weight,  and  we  find  out 
whether  it  is  such  a  multiple  only  after  the  atomic  weight  has  been 
determined.  The  atomic  weight  is  the  unit  weight  used  in  com- 
pounds, and  can  be  ascertained  only  by  a  study  of  compounds.  The 
molecular  weight  of  the  free  element  gives  us  only  a  value  which  we 
know  must  be  a  multiple  of  the  atomic  weight,  by  1  or  some  other 
integer.  Mol.  Wt.  =  At.  Wt.  X  x,  where  x  is  1  or  some  other  integer. 

Further  Discussion  of  the  Molecular  Formulae  of  Ele- 
mentary Substances.  —  Some  further  explanation  may  be  required, 
to  the  end  that  the  reader  may  be  reconciled  to  accepting  the  formulae 
C12,  O2,  and  so  forth.  In  the  first  place,  he  should  note  how  these 
formulae  arose.  If  we  accept  Avogadro's  hypothesis,  and  the  infer- 
ence from  it  to  the  effect  that  the  densities  of  gases  are  in  the  same 
ratio  as  the  weights  of  their  individual  molecules,  then  we  cannot 
escape  the  conclusion  to  which  measuring  the  relative  densities  of 
free  chlorine  and  hydrogen  chloride,  for  example,  leads.  The  ratio 
of  their  densities  is  70.9  :  36.45.  That  is  to  say,  the  relative  weights 
of  a  molecule  of  chlorine  and  a  molecule  of  hydrogen  chloride  stand 
in  this  ratio.  The  molecule  of  chlorine  is  nearly  twice  as  heavy  as 
the  molecule  of  the  compound,  and  there  cannot  therefore  be  a  whole 
molecule  of  chlorine  in  a  molecule  of  hydrogen  chloride.  In  fact,  we 
perceive  at  once  that  the  molecule  of  hydrogen  chloride  must  contain 
only  half  a  molecule  of  chlorine  (35.45) ,  together  with  half  a  molecule 
of  hydrogen  (1).  In  other  words,  if  the  molecule  of  free  chlorine 
were  to  be  taken  as  the  atom  of  the  element,  then  the  molecule  of 
hydrogen  chloride  would  contain  only  half  an  atom  of  chlorine, 
which  would  be  contrary  to  our  decision  to  take  as  atoms  quantities 
which  are  not  divided.  So  we  choose  the  other  horn  of  the  dilemma, 
and  say  that  the  specimen  of  chlorine  in  the  molecule  of  hydrogen 
chloride  is  a  whole  atom  and  that  therefore  the  amount  of  chlorine 
in  the  molecule  of  free  chlorine  is  two  atoms,  and  its  formula  C12. 
Similarly,  the  weight  of  hydrogen  in  the  molecule  of  hydrogen  chlo- 
ride is  1.008,  while  that  of  the  molecule  of  hydrogen  is  2.016,  so  that 


140  COLLEGE    CHEMISTRY 

there  are  two  atoms  in  the  molecule  of  free  hydrogen  and  its  formula 
is  H2.  Reasoning  in  like  manner  from  the  molecular  weights  of 
oxygen  (32)  and  water  (18)  we  reach  the  conclusion  that  the  mole- 
cule of  oxygen  is  diatomic  (02). 

Still  another  way  of  looking  at  the  same  facts  may  shed  light  on 
the  matter.  When  hydrogen  and  chlorine  combine,  one  volume  of 
each  of  these  gases  gives  two  volumes  of  hydrogen  chloride  (p.  122). 
Let  us  imagine  the  experiment  to  be  made  with  minute  volumes 
holding  one  hundred  molecules  each: 

HYDROGEN  CHLORIDE  HYDROGEN        CHLORINE 


100     100  came  from 


The  200  molecules  of  hydrogen  chloride  must  contain  at  least  200 
fragments  of  chlorine,  since  there  is  a  sample  in  each  molecule.  Now 
the  200  fragments  of  chlorine  came  from  a  volume  containing  only 
100  molecules  of  chlorine.  Each  of  these  must  therefore  have  been 
split  in  the  chemical  action.  Hence  the  molecules  of  free  hydrogen 
and  free  chlorine  contain  at  least  two  atoms.  If  we  consider  the 
molecular  formula  of  a  substance  as  representing  one  molecule  (see 
below),  the  equation  for  this  action  is: 

H2  +  C12  ->  2HCL 

There  are  two  molecules  on  each  side  of  the  equation,  and  this  corre- 
sponds with  the  fact  that  there  is  no  change  in  the  total  volume. 
Again,  we  find  that  one  volume  of  oxygen  furnishes  enough  of  the 
element  for  two  volumes  of  water  vapor  (p.  84).  We  infer  therefore 
that  each  molecule  of  oxygen  is  divided  into  two  parts  in  the  action. 
And  in  like  manner,  when  we  find  that  one  volume  of  phosphorus 
vapor,  in  combination  with  six  volumes  of  chlorine,  gives  four 
volumes  of  phosphorus  trichloride  vapor,  we  infer  that  every  mole- 
cule of  phosphorus  furnished  enough  of  the  element  for  four  molecules 
of  phosphorus  trichloride,  and  contained  therefore  four  atoms  (see 

Exercise  7). 

P4  +  6C12  -»  4PC13. 

The  simple  fact  that  hydrogen  and  oxygen,  when  mixed,  do  not 
combine  (p.  74)  may  assist  in  reconciling  us  to  the  diatomic  nature 
pf  their  molecules.  Some  part  of  the  mixture  has  to  be  heated 


APPLICATIONS   OF   MOLECULAR   EQUATIONS  141 

strongly  to  start  the  interaction.  Now  the  molecular  formulae,  H2 
and  02,  suggest  that  each  gas  is  really  in  combination  already  (with 
itself),  and  they  therefore  explain  to  some  extent  the  indifference  of 
the  gases  towards  one  another.  If  the  molecules  were  free  atoms, 
they  could  not  encounter  one  another  continually  as  they  move  about, 
and  yet  escape  combination  as  we  observe  that  they  do.  We  may 
imagine  that  the  primary  effect  of  heating  is  to  decompose  some  of 
the  molecules,  and  liberate  hydrogen  and  oxygen  in  the  atomic  con- 
dition, and  that  the  combination  of  these  atoms  starts  the  explosion 
of  the  whole  mass.  Of  course  this  explanation  is  based  upon  our 
hypothesis  and,  as  such,  is  of  the  same  imaginary  description  as 
everything  connected  with  that  hypothesis.  It  cannot  be  verified 
by  experiment. 

APPLICATIONS. 

Applications:  Interactions  Between  Gases.  —  According  to 
Avogadro's  hypothesis,  if  we  filled  a  succession  of  vessels  of  equal 
dimensions  with  different  gases,  and  could  arrest  the  motion  of  the 
particles  and  observe  their  disposition,  we  should  find  that  the  aver- 
age distance  from  particle  to  particle  would  be  the  same  in  all  cases. 
This  would  be  true  whether  our  vessels  were  filled  with  single  gases, 
with  homogeneous  mixtures,  or  with  gases  in  layers.  Such  being 
the  case,  if  any  chemical  change  is  brought  about  in  the  mass  which 
results  in  a  multiplication  of  the  molecules,  it  is  evident  that  the 
volume  will  have  to  increase  in  order  that  the  spacing  may  remain 
the  same  as  before.  If  any  chemical  action  results  in  a  diminution 
of  the  number  of  molecules,  then  a  shrinkage  must  take  place  in 
order  that  the  spacing  may  be  preserved  as  before.  Thus,  in  a  mix- 
ture of  hydrogen  and  chlorine,  according  to  our  hypothesis,  neigh- 
boring molecules  of  hydrogen  and  chlorine  simply  exchange  units, 
so  that  HH  4-  C1C1  becomes  HC1  +  C1H.  There  being  no  alteration 
in  the  number  of  particles,  no  change  in  volume  occurs.  In  the  case 
of  water,  on  the  other  hand, 

HH  +  OO  +  HH  becomes  HOH  +  HOH. 

Since  the  oxygen  molecules,  which  form  a  third  of  the  whole,  dis- 
appear into  the  molecules  of  hydrogen,  the  tendency  to  preserve 
spacing  results  in  a  diminution  of  the  volume  by  one-third  (p.  85). 
This  method  of  looking  upon  chemical  interactions  between  gases 
gives  us  the  nearest  sight  which  we  can  have  of  the  behavior  of  the 


142  COLLEGE   CHEMISTRY 

molecules  themselves.  We  cannot  perceive  the  individual  molecules, 
but,  in  consequence  of  the  spatial  arrangement  which  we  suppose 
them  to  observe,  the  change  in  the  whole  volume  of  a  large  aggre- 
gate of  molecules  enables  us  to  draw  conclusions  at  once  in  regard  to 
the  behavior  of  the  single  molecules  in  detail. 

Applications:  Molecular  Equations.  —  To  utilize  the  fore- 
going considerations,  chemists  always  employ  in  their  equations  the 
molecular  formulae  for  the  gases  and  easily  vaporized  substances  con- 
cerned. Thus  far,  we  have  used  the  equation: 

H     +     Cl    -»    HC1, 
WEIGHTS:  1.008        35.45        36.458 

and  the  information  it  contained  was  exhausted  when  we  had  placed 
below  the  symbols  the  weights  for  which  they  stood.  But  the 
molecular  equation  is  much  more  instructive.  The  following  shows 
the  interpretations  to  which  the  molecular  equation  is  subject: 

H2      +      C12      ->       2HC1. 

WEIGHTS:  2.016  g.         70.90  g.  2  X  36.458  (=  72.916)  g. 

VOLUMES:  22 . 4  1.  22 . 4  1.  2  X  22 . 4  1. 

MOLECULES:  11  2 

The  weights,  although  doubled,  show  the  same  proportions,  so 
that  questions  of  weight  are  answered  as  easily  as  before.  These 
weights,  however,  being  molecular  weights,  or  multiples  thereof,  can 
be  translated  at  once  into  volumes,  and  questions  about  volumes  can 
also  be  answered.  Finally  the  relative  numbers  of  each  kind  of 
molecules  can  be  read  from  this  equation,  for  the  coefficients  in 
front  of  the  formula  represent  these  numbers.  Where  no  coefficient 
is  written,  1  is  to  be  understood.  The  application  of  these  properties 
of  molecular  equations  is  illustrated  below.  Before  applying  these 
equations,  however,  we  must  first  learn  how  to  make  them. 

To  make  a  molecular  equation,  we  first  make  an  equation  according 
to  the  rules  already  explained  (p.  41).  An  equation  like  that  given 
for  the  interaction  of  oxygen  with  hydrogen  chloride  (Deacon's  pro- 
cess, p.  110):  2HC1  +  O->H2O  +  2C1,  is  the  result.  Then  we 
adjust  the  equation  so  that  molecular  formulae  are  used  throughout. 
2C1  becomes  at  once  C12.  The  oxygen,  however,  must  also  appear 
as  O2,  or  a  multiple  of  this,  in  such  equations.  Hence  the  whole 
equation  must  be  multiplied  by  2: 

4HC1  +  O2  -»  2H2O  +  2CL,. 


APPLICATIONS   OF   MOLECULAR   EQUATIONS  143 

Again,  the  equation  for  the  preparation  of  chlorine  from  potassium 
permanganate  (p.  Ill)  becomes: 

2KMn04  +  16HC1  ->  8H2O  +  2KC1  +  2MnCl2  +  5C12. 

Every  equation  containing  an  odd  number  of  atoms  of  a  substance 
whose  molecules  are  diatomic  must  be  multiplied  by  2.  Again,  mer- 
curic oxide  decomposes  to  give  mercury  vapor  and  oxygen  (p.  43), 
and  the  molecules  of  mercury  are  monatomic  and  those  of  oxygen 
diatomic,  so  we  write: 

2HgO  ->  2Hg  +  02. 

Finally,  the  formulae  of  substances  which  are  solid  or  liquid,  and 
cannot  be  easily  vaporized,  are  written  in  the  simplest  terms.  Thus, 
since  substances  like  the  copper  in  the  following  equation  are  invola- 
tile,  the  molecular  weights  of  such  substances  are  unknown,  and  their 
molecular  formulae  likewise:  2Cu  +  O2  — >  2CuO.  Furthermore,  in 
the  case  of  substances  which  can  be  volatilized,  although  the  molec- 
ular weights  and  molecular  formulae  may  therefore  be  known,  we 
do  not  usually  employ  the  molecular  formula?  if  the  substance  is  not 
used  in  the  form  of  vapor  in  the  laboratory.  Thus,  the  molecular 
formula  of  phosphoric  anhydride  is  P4OIO  (p.  132) .  But  we  generally 
make,  and  use,  only  the  solid  form,  and  not  the  vapor,  in  actual  work. 
Hence  the  action  with  water  is  usually  written  as  we  have  given  it 
(p.  51),  rather  than  in  the  form:  P4O10  +  6H2O  ->  4H3PO4. 

The  applications  of  the  properties  of  molecular  equations  (which  will 
be  used  exclusively  hereafter)  may  now  be  illustrated  in  detail. 

Applications:  To  Arithmetical  Problems.  —  1.  When  a 
problem  in  regard  to  weights  of  material  used  or  produced  in  a 
given  action  is  to  be  solved,  the  molecular  equation  is  to  be  written 
and  the  weights  inserted  beneath  the  formulae.  The  mode  of  calcu- 
lation has  been  described  already  (p.  43). 

2.  When  a  problem  involving  weights  and  volumes  is  to  be  solved, 
the  molecular  equation  is  to  be  written,  and  both  the  weights  and 
volumes  are  to  be  inserted.  Note,  however,  that  only  the  volumes 
of  the  substances  in  the  gaseous  condition  are  considered. 

For  example,  what  volume  of  oxygen  is  obtained  from  60  g.  of 
potassium  chlorate?  The  molecular  equation,  made  as  shown  above 
(p.  142),  together  with  the  full  interpretation,  are  as  follows: 


144  COLLEGE,  CHEMISTRY 


f 

:  •<  * 
t 


2KC103  ->  2KC1 

2  (39.15  +  35.45  +  48)     2(39.15  +  35.45) 


WEIGHTS 

245.  2  g.  149.  2  g.  96  g. 

VOLUMES:  3  X  22,41. 

(Observe  that  no  volumes  are  given  under  the  chlorate  and  chloride 
of  potassium.  This  is  because  their  volumes  in  the  gaseous  condition 
can  be  of  no  practical  use,  since  they  are  solids  which  are  melted, 
but  not  vaporized  during  this,  or  any  action  in  which  we  employ  them.) 
Now,  as  to  the  problem  in  hand,  it  is  concerned  with  a  weight  of 
potassium  chlorate  and  a  volume  of  oxygen.  Reading  from  the 
equation,  our  information  on  these  points  is  that  245.2  g.  of  potassium 
chlorate  give  67.2  liters  (observe  that  the  coefficients  are  used,  as 
well  as  the  molecular  weights,  in  these  numbers)  of  oxygen  at  0°  and 
760  mm.,  and  the  question  is,  What  volume  will  60  g.  give?  By 
proportion,  245.2  g.  :  67.2  1.  :  :  60  g.  :x  1.,  where  x  =  16.4  liters.  If 
a  different  temperature  and  pressure  had  been  specified,  either  the 
volume  in  the  equation,  or  the  answer,  would  have  had  to  be  con- 
verted by  rule  to  the  given  conditions.  It  saves  time  not  to  write 
out,  as  above,  the  whole  interpretation,  but  only  the  parts  required. 
For  example,  if  the  question  is  :  What  volume  of  chlorine  is  needed 
to  give  25  g.  of  aluminium  chloride,  we  may,  if  we  choose,  omit  all 
the  data  excepting  the  volume  of  the  chlorine  and  the  weight  of  the 
aluminium  chloride,  thus: 

2A1     +     3C12      ->      2A1C13 

3X22.41.     2  X  133.  45  g. 

The  volume  of  chlorine  required  is  25  X  3  X  22.4  H-  (2  x  133.45)  liters. 
These  illustrations  show  the  method  of  calculating  actual  volumes 
(see  Exercises  8,  9). 

3.  If  the  question  concerns  relative  volumes  only,  then  it  is  simplest 
to  use  the  interpretation  of  the  equation  in  terms  of  molecules.  For 
example,  What  relative  volumes  of  hydrogen  chloride  and  oxygen 
are  required  in  Deacon's  process?  The  molecular  equation  is  (p.  142)  : 

4HC1  +  O2  ->  2H20  +  2C12. 
MOLECULES:  41  22 

Since  equal  numbers  of  molecules  of  gases  occupy  equal  volumes, 
the  proportion  4  molecules  of  hydrogen  chloride  to  1  molecule  of 
oxygen  shows  the  ratio  to  be  4  :  1  by  volume.  Similarly,  every  4 


APPLICATIONS   OF    MOLECULAR   EQUATIONS  145 

molecules  of  hydrogen  chloride  give  2  molecules  of  chlorine,  so  that 
the  ratio  of  these  substances  by  volume  is  4  :  2,  or  2  :  1. 

In  regard  to  the  water,  since  that  is  not  a  gas  at  common  tem- 
peratures, the  question,  if  asked,  must  be  more  specific:  What  are 
the  relative  volumes  of  steam  and  chlorine  in  the  product,  as  com- 
monly delivered  by  the  action  at  400°?  It  is  2  :  2,  or  1  :  1.  What 
are  the  relative  volumes  of  water  and  chlorine,  after  the  products 
have  cooled  to  room  temperature?  The  water  is  no  longer  a  gas, 
so  that  it  occupies,  relatively,  almost  no  volume.* 

What  is  the  total  volume-change  in  the  foregoing  action  above 
100°?  It  is  a  change  from  5  molecules  to  4.  The  volume  changes 
in  the  same  ratio.  But  at  0°  the  volume-change  is  from  5  volumes 
to  2,  for  the  water  does  not  appreciably  add  to  the  volume  of  the 
products  (see  Exercises  10,  12). 

4.  When  we  know  the  molecular  formulae  of  the  single  substances 
concerned  in  an  action,  the  equation  can  be  made,  and  the  relative 
volumes  determined,  without  actual  measurement.  For  example:  What 
volume-change  will  be  observed  when  a  mixture  of  carbon  monoxide 
and  oxygen  has  exploded,  and  the  temperature  has  once  more 
reached  that  of  the  room?  The  molecular  formulae  are  CO,  O2,  and 
C02.  The  equation  representing  the  weights  is  CO  +  0  — >  C02. 
The  molecule  of  oxygen,  however,  being  02,  we  cannot  employ  less 
than  this  quantity  in  a  molecular  equation,  so  that  the  equation 
becomes:  2CO  +  02 ->  2CO2. 

Three  molecules,  therefore,  give  two,  throughout  the  whole  mass, 
and  therefore  three  volumes  will  become  two,  if  the  pressure  and 
temperature  are  the  same  at  the  beginning  and  end  of  the  action. 

If  we  remember  that  all  volatile  compounds  of  carbon  and  hydro- 
gen burn  to  form  water  and  carbon  dioxide,  the  molecular  equation 
for  any  such  combustion  may  easily  be  made,  and  the  volumes  of  all 
the  materials  ascertained.  When  water  is  a  product,  only  its  volume 
as  steam  is  given  by  the  equation  (see  Exercises  11,  12). 

*  Of  course  if  an  exact  answer  must  be  given,  it  can  be  given.  But  for 
this  we  require  the  weight  and  specific  gravity  of  the  product.  Thus,  21T,O 
represents  2  X  18  g.  of  water.  The  sp.  gr.  of  water  is  1.  Therefore  the 
volume  of  water  formed  is  36  c.c.  The  volume  of  2C12  is  2  X  22.4,  or  44.8 
liters  at  0°.  The  ratio  of  water  to  chlorine  by  volume  at  0°  is  therefore  36  : 
44,800.  But,  as  a  rule,  we  simply  give  the  volumes  of  solids  and  liquids  as 
zero,  compared  with  those  of  the  gases  concerned  in  the  same  action. 


146  COLLEGE   CHEMISTRY 

5.  Knowing  by  heart  the  molecular  formulae  of  gaseous  substances, 
as  we  must  know  them  for  many  purposes,  it  is  unnecessary  to 
burden  our  minds  with  other  data  in  regard  to  the  relative  weights  of 
gases.  Is  hydrogen  chloride  (HC1)  heavier  or  lighter  than  carbon 
dioxide  (CO2)?  These  formulae  represent  the  weights  of  equal 
volumes  (22.4  L),  namely  36.45  g.  and  44  g.,  respectively.  Hence 
the  former  gas  is  a  little  lighter.  Remembering  that  the  G.M.V.  of 
air  weighs  28.955  g.,  we  can  compare  the  weight  of  any  gas  with  that 
of  air  in  the  same  way. 

What  are  the  relative  weights  of  acetylene  (C2H2,  p.  132)  and 
sulphur  dioxide  (SO2)  as  compared  with  air?  The  G.M.V.  cube 
holds  formula-weights  of  the  first  two,  namely  26  g.  and  64  g.,  and 
28.955  g.  of  air.  Hence  acetylene  is  a  little  lighter  than  air,  and 
sulphur  dioxide  more  than  twice  as  heavy  (see  Exercise  13). 

Applications :  To  Cases  of  Dissociation.  —  Several  substances 
yield  smaller  values  for  their  densities,  and  therefore  molecular 
weights,  when  the  densities  are  measured  at  higher  temperatures. 
This  indicates  that  the  molecules  have  become  lighter,  and  can  only 
mean  that  decomposition  has  taken  place  in  consequence  of  the 
heating.  Behavior  of  this  kind  is  shown  both  by  compounds  and 
by  simple  substances. 

For  example,  phosphorus  pentachloride  PC15,  although  a  solid, 
can  be  converted  into  vapor  without  much  difficulty.  Its  molecular 
weight,  if  it  underwent  no  chemical  change  during  the  volatilization, 
would  be  31  +  177.25  =  208.25.  The  density  actually  observed  at 
300°  gives  by  calculation  not  much  more  than  half  this  value.  The 
direct  inference  from  this  is,  that  the  molecules  have  only  half  the 
(average)  weight  that  we  expected :  or,  in  other  words,  are  twice  as 
numerous  as  we  expected.  The  explanation  is  found  when  we 
examine  the  nature  of  the  vapor  more  closely.  We  find  that  it  is  a 
mixture  of  phosphorus  trichloride  and  free  chlorine,  resulting  from 
a  chemical  change  according  to  the  equation :  PC15  <=±  PC13  +  C12. 
The  low  value  of  the  density  thus  tells  us  that  dissociation  has 
taken  place.  From  the  value  of  the  density  at  various  tempera- 
tures, we  may  even  calculate  the  proportion  of  the  whole  material 
which  is  dissociated.  At  300°  it  is  97  per  cent;  at  250°,  80  per  cent; 
and  at  200°,  48.5  per  cent.  Thus,  when  the  temperature  is  lowered, 
progressive  recombination  takes  place  and  the  proportion  dissociated 


APPLICATIONS    OF   MOLECULAR    EQUATIONS  147 

becomes  less.     Finally  the  vapor  condenses  and  yields  the  original 
solid. 

Again,  sulphur  boils  at  445°,  but  can  be  vaporized  at  a  temperature 
as  low  as  193°,  under  very  low  pressure.  At  this  temperature  the 
density  of  the  vapor  gives  the  molecular  weight  256  (=  8X32),  and 
the  molecular  formula  S8.  That  is  to  say,  the  G.M.V.  holds  256  g. 
of  the  vapor  at  193°.  At  800°,  however,  the  density  is  only  one- 
fourth  as  great,  and  the  G.M.V.  holds  only  64  g.  (S2).  This  means 
that  256  g.  now  occupy  four  times  as  large  a  volume  as  before,  and 
the  increase  is  additional  to  the  effect  of  the  mere  thermal  expansion, 
which  is  allowed  for  in  the  calculation  and  eliminated.  Hence  the 
molecules  have  dissociated.  At  1700°  the  molecular  formula  is  still 
S2,  so  that  this  represents  the  limit  of  dissociation :  S8  +±  4S2.  When 
the  vapor  is  cooled,  the  density  increases  once  more  and  at  193° 
recovers  completely  the  greater  value.  Similar  observations  show 
that  phosphorus  vapor  at  313°  is  all  P4,  but  at  1700°  half  of  it  has 
dissociated  into  P2.  Iodine  vapor,  up  to  445°,  is  all  I2.  Beyond 
this  temperature  the  density  diminishes,  and  when  1700°  is  reached 
the  vapor  is  all  I.  Thus  the  molecules  are  diatomic  at  low  tempera- 
tures and  monatomic  at  high  ones.  The  densities  of  oxygen,  hydro- 
gen, and  chlorine  are  not  measurably  affected  by  heating  to  1700°, 
so  that  their  diatomic  molecules  exist  from  temperatures  far  below 
0°  up  to  1700°,  and  are  evidently  very  stable. 

Applications :  Finding  the  Atomic  Weight  of  a  New  Ele- 
ment.— By  way  of  reviewing  the  principles  explained  in  the  chapter, 
let  us  apply  them  to  the  imaginary  case  of  a  newly  discovered  element. 
The  bromide  of  the  element  is  found  to  be  easy  of  preparation  and 
to  be  volatile.  The  bromide  contains  30  per  cent  of  the  element, 
and  its  vapor  density  referred  to  air  is  11.8.  The  analysis  can 
always  be  made  much  more  accurately  than  the  measurement  of 
vapor  density,  so  that  the  former  number  is  more  trustworthy  than 
the  latter. 

To  find  the  equivalent  of  the  element,  that  is,  the  amount  com- 
bined with  80  parts  (the  equivalent)  of  bromine,  we  have  the  pro- 
portion 70  :  30  : :  80  :  x,  from  which  x  =  34.3.  The  atomic  weight 
must  be  this,  or  some  small  multiple  of  it. 

The  G.M.V.  of  air  weighs  28.955  g.  (p.  126).  Hence  the  same 
volume  of  the  vapor  of  this  bromide,  which  is  11.8  times  as  heavy  as 


148  COLLEGE   CHEMISTRY 

air,  will  weigh  28.955  X  11.8,  or  341.67  g.  This  is  therefore  the 
molar  weight  of  the  compound. 

Now  30  per  cent  of  this  is  the  new  element: 
341.67  X  30  -  100  =  102.5. 

Three  times  the  equivalent  weight  is  the  multiple  nearest  to  this 
number,  3  X  34.3  =  102.9,  the  difference  being  due  to  error  in 
determining  the  density.  So  long  as  no  other  volatile  compound 
is  known,  we  adopt  this  as  the  atomic  weight.  The  rest  of 
the  molar  weight  (240  parts)  is  bromine.  Thus  the  formula  of 
the  compound  is  ElBr3,  and  from  this  we  see  that  the  element  is 
trivalent. 

In  case  no  volatile  compound  of  the  element  can  be  formed,  the 
equivalent  is  measured  as  before.  Then  some  of  the  free  simple 
substance  is  made,  say  by  electrolysis,  and  its  specific  heat  is  deter- 
mined. The  sp.  ht.  is  about  0.063.  Application  of  Dulong  and 
Petit's  law  then  gives  the  atomic  weight.  The  product  34.3X0.063 
is  equal  to  2.161.  Hence,  the  equivalent  must  be  multiplied  by  3 
to  give  the  atomic  weight,  for  this  raises  the  product  to  6.48,  which 
is  within  the  limits.  Thus  the  value  of  the  atomic  weight  is  102.9, 
as  before. 

Replies  to  Questions  about  Difficulties.  —  The  beginner 
always  becomes  confused  over  one  or  more  of  the  points  raised  by 
the  following  questions: 

Why  was  32  g.  of  oxygen  taken  as  the  standard  for  molecular 
weights,  rather  than  16  g.?  Read  p.  132  and  footnote  to  p.  131. 

If  O2  is  the  smallest  mass  of  oxygen,  why  do  we  have  formulae  like 
H2O  and  HC1O?  O2  is  the  smallest  mass  of  free  oxygen,  but  in  com- 
bination half  as  much  occurs  in  many  molecules.  Read  pp.  132,  137, 
and  138. 

Why  is  not  the  atomic  weight  of  art  element  ascertained  by  simply 
measuring  the  density  of  the  elementary  substance?  Read  pp.  139 
and  147. 

Can  we  not  deduce  the  valence  of  an  element  from  knowing  the 
number  of  atoms  in  its  molecules,  and  vice  versa  f  Some  molecular 
formulae  and  valences  are:  H2T,  O2n,  CL/,  Zn",  also  Hg  (univalent 
and  bivalent),  P4  (trivalent  and  quinquivalent)  and  S8  (bivalent 
and  sexivalent).  There  is  no  relation,  either  observable  or  to  be 
expected. 


. 

APPLICATIONS   OF   MOLECULAR   EQUATIONS  149 

Do  the  molecular  weights,  oxygen  =  32  and  hydrogen  =  2  mean 
that  the  molecules  of  oxygen  are  larger  than  are  those  of  hydrogen? 
This  is  the  ratio  of  their  weights,  but  none  of  the  phenomena  dis- 
cussed in  this  chapter  are  influenced  appreciably  by  their  relative 
sizes,  and  therefore  none  of  them  give  any  information  on  the 
subject.  Read  the  footnote  to  p.  129. 

Exercises.  —  1.  The  weight  of  1 1.  of  a  gas  at  0°  and  760  mm.  is 
5.236  g.  What  is  the  density  referred  to  air  and  to  hydrogen,  and 
what  is  the  molecular  weight  (pp.  126,  129)? 

2.  The  density  of  a  gas,  referred  to  air,  is  6.7.     What  is  the  weight 
of  1 1.  (p.  126),  and  what  is  the  molecular  weight  (p.  147)? 

3.  The  molecular  weight  of  a  substance  is  65.     What  is  the  density 
referred  to  air,  and  what  is  the  weight  of  1  1.? 

4.  The  chloride  of  a  new  element  contains  38.11  per  cent  of  chlo- 
''  rine  and  61.89  per  cent  of  the  element.     The  vapor  density  of  the 

compound  referred  to  air  is  12.85.  What  is  the  atomic  weight  of 
the  element,  so  far  as  investigation  of  this  one  substance  can  give 
it  (p.  130)?  What  is  its  valence? 

5.  If  the  molecular  weight  of  oxygen  were  taken  as  100,  what 
would  be  the  volume  of  the  G.M.V.  (p.  126)?    What  would  be  the 
molecular  weight  of  water,  and  what  would  be  the  atomic  weights 
of  hydrogen  and  chlorine  (pp.  126,  132)? 

6.  In  future  nothing  but  molecular  formulae  of  fre/e  elements  must 
be  used   (p.  138).     Write  in  molecular  form  ten  of  the  equations 
involving  gases  which  are  found  in  the  preceding  chapters. 

7.  If  a  new  form  of  oxygen  were  found,  such  that  one  volume  of 
it  required  four  volumes  of  hydrogen  to  produce  water,  what  would 
be  its  molecular  formula  (p.  140)?     What  would  be  the  weight  of 
22.41.? 

•  8.   What  volume  of  oxygen  at  10°  and  750  mm.  is  obtainable  by 
heating  50  g.  of  barium  peroxide  (pp.  46,  143-144)? 
.  ,    9.   What  volume  of  oxygen  at  20°  and  760  mm.  is  required  to 
Vconvert  16  g.  of  iron  into  dehydrated  rust  (Fe2O3)  (p.  144)? 

10.  Write  out  the  molecular  equations  for  the  interactions  of 
methane  and  chlorine  (pp.  115,  142);  and  for  the  burning  of  phos- 
phorus (vapor)  in  oxygen  (pp.  132,  138,  145).  Deduce  the  volume 
relations  of  the  initial  substances,  and  of  the  products,  at  various 
temperatures  in  each  case. 

3  ^  -*     ?    ^.. 

*>•;•)  •••"  siv  X 


150  COLLEGE   CHEMISTRY 

11.  Write  out  the  molecular  equations  for  the  interactions  of 
acetylene  and  oxygen  (p.  145),  and  of  alcohol  vapor  (b.-p.  78°)  and 
oxygen.     Deduce  the  volume  relations  of  the  initial  substances  and 
of  the  products  at  0°  and  at  100°  in  each  case. 

12.  The  molecular  weight  of  cyanogen  is  52.08.     What  is  its  den- 
sity referred  to  air,  and  what  the  weight  of  1  1.  at  0°  and  760  mm.? 
It  contains  46.08  per  cent  carbon  and  53.92  per  cent  nitrogen.     What 
is  the  formula  of  the  substance  (p.  41)?    Exploded  with  oxygen  it 
forms  carbon  dioxide  and  free  nitrogen.     What  will  be  the  relative 
volumes  of  the  materials  before  and  after  the  interaction  (p.  145)? 

13.  What  are  the  relative  weights  of  equal  volumes  of  hydrogen 
sulphide  (H2S),and  hydrogen  iodide  (HI),  compared  with  air  (p.  145)? 


I 


CHAPTER  XIII 
THE   ATOMIC   HYPOTHESIS 


To  determine  the  nature  of  chemical  phenomena  requires,  as  we 
have  found,  very  elaborate  experimentation.  And  this  has  to  be 
followed  by  still  more  elaborate  reasoning  before  a  systematic  state- 
ment of  the  precise  nature  of  the  change  can  be  made.  Yet,  when 
all  this  has  been  done,  we  are  still  unable  to  form  a  clear  conception 
of  what  manner  of  procedure  the  change  follows,  for  the  details  are 
entirely  inaccessible  to  observation.  We  should  like  to  know  pre- 
cisely how  chemical  union  is  consummated,  and  how  chemical 
exchange  is  carried  out.  We  should  like  to  account  for  the  fact  that 
the  automatically  adjusted  proportions  of  the  materials  used  in  every 
chemical  change  are  entirely  uninfluenced  by  temperature  and  other 
conditions.  Above  all,  we  should  like  to  know,  if  possible,  what 
state  of  affairs  determines  the  employment  of  an  individual  unit 
weight  by  each  element  in  all  its  combinations.  None  of  these 
questions  can  be  answered,  however,  because  nothing  can  be  seen 
which  suggests  any  answer. 

As  usual  in  cases  of  this  kind  we  construct  an  imaginary  mech- 
anism, a  formulative  hypothesis  (p.  86)  to  account  for  the  facts.  In 
doing  so,  we  are  not  under  the  illusion  that  we  are  discovering  the 
actual  machinery.  We  realize  that  we  are  simply  making  a  sort  of 
diagram  which  will  assist  our  thought  about  the  thing  itself.  Now 
a  slight  addition  to  the  molecular  hypothesis  readily  furnishes  pre- 
cisely what  we  need. 

Atomic  Hypothesis.  —  According  to  the  molecular  hypothesis,  all 
matter  is  made  up  of  small  discrete  particles,  each  of  which  has  the 
same  composition  as  has  the  body  as  a  whole.  It  is  difficult,  there- 
fore, to  avoid  the  conception  that  the  different  materials  in  each 
compound  molecule  are  more  or  less  distinct  entities  also.  Hence 
we  make  this  the  basis  of  a  new  hypothesis,  and  attribute  to  these 
constituent  parts  of  molecules  the  properties  of  the  chemical  unit 

151 


152  COLLEGE    CHEMISTRY 

weights  which  are  closely  related  to  them.  Thus,  these  parts  must 
move  from  one  state  of  combination  to  another  without  alteration  in 
their  mass.  Each  kind  must  be  composed  of  a  distinct  variety  of 
matter. 

These  parts  of  molecules  which  we  thus  suppose  to  be  permanent, 
coherent  masses,  are  named  atoms.  This  word  signifies  bodies  which 
are  not  disintegrated  (Gk.  aro/^o?,  uncut,  i.e.  not  cut).  The  relative 
weights  of  these  imaginary  masses  are  the  atomic  weights. 

It  should  be  noted  that  although  atoms,  like  molecules,  are  fic- 
tions, the  atomic  weights,  since  they  are  measured  by  experimental 
methods,  are  real.  They  originally  received  this  name  when  Dalton, 
an  English  schoolmaster  of  Manchester,  succeeded  in  unraveling  the 
complications  of  the  chemical  composition  of  substances  by  the  help 
of  this  very  hypothesis,  and  realized  for  the  first  time  (1802)  that 
the  possession  by  all  elements  of  individual  chemical  unit  weights 
lay  at  the  basis  of  the  whole  system. 

We  may  sum  up  all  that  the  facts  require  us  to  assume  about  atoms 
by  saying:  Atoms  are  the  units  of  which  molecules  are  aggregates. 
Those  of  like  kind  have  equal  masses,  and  differ  from  those  of  other 
kinds  both  in  mass  and  the  kind  of  material  of  which  they  are  made. 
The  fundamentally  different  kinds  of  materials  are  the  chemical 
elements. 

We  must  carefully  avoid,  for  the  present,  the  introduction  of 
unnecessary  complications  by  inventing  for  atoms  any  other  prop- 
erties, such  as  size,  shape,  or  color,  gratuitously.  It  is  to  be  noted, 
particularly,  that  there  are  no  facts  in  chemistry  which  require  us 
to  suppose  that  atoms  are  incapable  of  disintegration.  The  molecule 
of  oxygen,  for  example,  might  be  a  cluster  of  particles.  Certain 
facts  (p.  140)  require  us  to  suppose  that  this  cluster  is  divisible  into 
two  parts,  but  for  chemical  purposes  the  half  cluster,  with  relative 
weight  16,  is  as  small  a  subdivision  as  we  need.  So  we  take  the  half 
molecule  as  the  chemical  unit  quantity,  without  inquiring  further 
into  its  structure.  It  would  only  occasion  the  risk  of  conflict  with 
facts  still  to  be  discovered  if  we  elaborated  our  fiction  any  further 
than  is  absolutely  necessary.  Indeed,  the  phenomena  of  radio- 
activity have  already  compelled  us  to  suppose  that  there  are  particles 
much  smaller  than  atoms  (see  Radium).  Without  this  assumption 
the  new  facts  cannot  be  accounted  for  in  harmony  with  the  molecular 
hypothesis. 


ATOMIC    HYPOTHESIS  153 

The  Atomic  Hypothesis  and  Chemical  Change.  —  The  change 
from  chemical  units  to  atoms  is  so  slight  that  the  application  of  this 
hypothesis  to  the  description  of  chemical  phenomena  is  very  readily 
made.  The  description  gains  in  concreteness,  however,  from  the 
change.  Thus,  we  consider  iron  from  every  source  to  be  made  of 
minute  portions  of  iron  matter  which  are  all  alike  in  weight,  and  pre- 
sumably also  in  their  other  properties.  Similarly,  all  specimens  of 
sulphur  are  made  of  minute  particles  of  sulphur  having  exactly  the 
same  weight.  The  weight  of  a  sulphur  atom,  however,  is  different 
from  that  of  an  iron  atom.  When  visible  portions  of  the  two  sub- 
stances unite,  we  conceive  the  operation  to  consist  in  the  union  of 
each  atom  of  iron  with  one  atom  of  sulphur  to  produce  a  molecule  of 
ferrous  sulphide.  The  consummation  of  this  union  between  multi- 
tudes of  pairs  of  the  respective  kinds  of  atoms  in  every  second  of 
time  results  in  a  chemical  transformation  whose  progress  is  percep- 
tible by  the  senses. 

In  more  complex  chemical  changes,  a  further  correspondence 
between  the  hypothesis  and  the  facts  comes  to  our  notice.  For 
example,  when  sulphuric  acid  acts  upon  sodium  chloride  to  produce 
hydrogen  chloride: 

NaCl  +  H2S04  -»  HC1  +  NaHSO4, 

one  part  by  weight  of  hydrogen  takes  the  place  of  23  parts  by  weight 
of  sodium,  and  combines  with  35.45  parts  of  chlorine  to  form  the 
hydrogen  chloride.  This  is  the  way  we  state  the  change  when  we 
refer  to  measurable  quantities.  According  to  this  hypothesis,  the 
operation  consists  in  the  repetition,  millions  of  times  over  within  a 
small  amount  of  material,  of  the  substitution  of  one  atom  of  hydrogen 
for  one  atom  of  sodium  to  form  a  molecule  of  hydrogen  chloride. 
The  special  fact  which  we  notice  is  that  the  atom  of  hydrogen  suffices 
exactly  to  occupy  the  place  of  the  atom  of  sodium.  If  it  were  some- 
what too  heavy  or  somewhat  too  light,  then  a  single  unit  weight  for 
hydrogen  would  not  describe  the  proportions  of  the  element  in  all  its 
combinations,  a  situation  which  would  be  contrary  to  the  fact 
recorded  in  the  law  of  combining  weights  (p.  34).  The  remarkable 
fact  about  this,  and  all  other  double  decompositions  of  the  same 
sort,  is  that  the  little  masses  of  the  various  elements  exchange  places, 
without  any  alterations  to  fit  the  new  compound  being  required. 
Hence  our  assumption  that  the  atoms  are  permanent,  coherent 


154  COLLEGE   CHEMISTRY 

wholes.  This,  of  course,  is  simply  stating  in  terms  of  the  atomic 
hypothesis  the  facts  which  underlie  the  conception  of  unit  weights. 
A  chemical  phenomenon  as  we  observe  it,  then,  is  imagined  to  consist  in 
some  systematic  liberation,  combination,  or  exchange  of  atoms,  according 
to  a  definite  scheme,  and  repeated  many  millions  of  times  (once  with 
each  molecule)  in  a  body  or  mixture  of  bodies. 

TJie  Atomic  Hypothesis  and  the  Quantitative  Laws.  — The 

idea  of  atoms  simply  crystallizes  somewhat  more  definitely  the  con- 
ception of  chemical  unit  weights.  Hence,  it  follows  of  necessity  that 
the  quantitative  laws  of  chemical  combination,  out  of  which  the 
latter  arose,  will  be -found  to  be  entirely  in  harmony  with  the  atomic 
hypothesis.  The  definite  composition  of  each  compound,  for  exam- 
ple, corresponds  to  the  hypothesis  that  each  substance  is  made  up  of 
a  specific  kind  of  molecules,  all  of  which,  in  turn,  contain  the  same 
kind  and  number  of  atoms. 

The  conception  of  valence  (p.  70)  suggests,  in  terms  of  this 
hypothesis,  that  some  atoms  unite  with  but  one  other  atom,  habitu- 
ally (NaCl).  Some,  however,  unite  with  two  of  the  first  kind  (ZnCl2), 
or  with  one  other  of  their  own  kind  (ZnO),  still  others  with  three 
atoms  of  the  first  kind  (A1C13),  and  so  forth.  In  other  words,  it 
involves  the  assumption  that  each  kind  of  atom  has  a  limited  capa- 
city for  holding  other  atoms  in  combination.  Thus,  taking  the  most 
crudely  mechanical  view  of  the  matter,  we  might  elaborate  the 
hypothesis  by  suggesting  that  there  is  a  limit  to  the  number  of  points 
at  which  atoms  may  be  attached  to  one  another.  When  one  atom 
of  chlorine  is  attached  to  one  of  sodium,  the  combining  capacity  of 
each  is  exhausted.  When  one  atom  of  hydrogen  is  attached  to  one 
atom  of  oxygen,  one  combining  capacity  of  the  oxygen  still  remains 
(H— O— ),  and  can  be  satisfied  by  one  more  atom  of  hydrogen  or  of 
some  other  element. 

Equations  and  the  Atomic  Hypothesis.  —  Since  equations  are 
simply  records  of  the  kinds  and  quantities  of  matter  taking  part  in 
chemical  changes,  they  require  no  addition  to  the  atomic  hypothesis. 
The  symbols,  which  originally  represented  unit  weights  of  the  various 
elements,  may  stand  equally  well  for  atoms.  Hence,  in  the  current 
language  of  chemistry,  potassium  chlorate  (KC1O3)  is  composed  of 
one  atom  each  of  potassium  and  of  chlorine,  and  three  atoms 


ATOMIC   HYPOTHESIS  155 

of  oxygen,  in  each  molecule.     The  word  "  atom  "  stands  for  unit 
weight,  and  the  word  "  molecule  "  for  molecular  weight. 

Concretely,  when  these  terms  are  used,  the  chemical  equation 
represents  one  minute  specimen  of  the  change  which  is  taking  place 
throughout  the  whole  mass.  It  shows  enough  molecules  and  atoms 
to  furnish  a  complete  example  of  what,  repeated  millions  of  times, 
will  constitute  a  chemical  change  in  a  visible  amount  of  material. 
There  are  a  few  details  which  this  point  of  view  suggests.  For 
example,  the  equation  HgO  — »  Hg  +  O  shows  the  weights,  but  does 
not  include  a  complete  sample  of  the  materials  in  the  point  of  view 
of  this  hypothesis.  The  smallest  fragment  of  free  oxygen  which  we 
can  obtain  is  made  of  two  atoms.  Hence  two  molecules  of  mercuric 
oxide  are  required  to  furnish  them.  Thus  the  minimum  number 
of  molecules  which  would  suffice  for  carrying  out  this  change  on  a 
molecular  scale  is  2HgO  — >  2Hg  +  O2.  Similarly,  since  the  small- 
est discrete  portions  of  free  hydrogen  and  chlorine  contain  two  atoms 
each,  we  must  write  the  equation  H2  +  C12  — >  2HC1  (p.  140). 

May  other  Properties,  Besides  Definite  Mass  and  Specific 
Material,  be  attributed  to  an  Atom?  —  In  one  direction,  namely, 
that  of  accounting  for  the  properties  of  compounds,  no  attempt  has 
been  made  to  adapt  the  atomic  hypothesis  so  as  to  explain  the  facts. 
Thus,  two  hydrogen  atoms  in  a  molecule  give  a  gas  almost  insoluble 
in  water;  two  chlorine  atoms,  a  gas  which  is  moderately  soluble;  but 
one  of  each  gives  hydrogen  chloride,  which  dissolves  in  water  in 
extraordinary  quantities.  So  also,  colorless  substances  give,  by 
chemical  union,  strongly  colored  ones,  and  odorless  substances,  by 
chemical  union,  strongly  odorous  ones.  Putting  pieces  of  iron  and 
sulphur  side  by  side  causes  absolutely  no  change  in  the  properties  of 
either.  And  yet  this  hypothesis  compels  us  to  assume  that  if  the 
particles  are  made  fine  enough,  and  placed  close  enough  to  one 
another,  the  individual  properties -of  the  constituents  will  entirely 
disappear.  Hitherto  we  have  failed  to  think  of  any  qualities  which 
might  be  attributed  to  the  atoms  in  order  to  account  for  facts  of 
this  class.  Why  should  oxygen  (O2)  and  ozone  (03)  be  so  different 
in  behavior,  although  the  atomic  theory  hints  at  nothing  but  a  sub- 
stitution of  three  atoms  for  two?  What  atomic  properties  shall 
account  for  the  difference  between  red  and  yellow  phosphorus? 
Usually,  in  explaining  these  matters,  we  forsake  atoms  and  employ 


156  COLLEGE   CHEMISTRY 

the  conception  of  energy.  We  say  that  ozone  has  more  energy  than 
oxygen,  and  yellow  phosphorus  more  than  red  phosphorus. 

There  are  certain  facts,  however,  such  as  the  chemical  actions 
known  as  internal  rearrangements  (p.  10),  which  are  successfully 
explained  by  supposing  that  the  atoms  composing  a  given  kind  of 
molecule  are  disposed  towards  one  another  in  a  definite  geometrical 
arrangement.  We  speak  then  of  the  structure  of  the  molecule  and 
the  constitution  of  the  substance,  and,  to  represent  our  conclusions, 
we  use  graphic  formulae  (q.v.). 

Summing  this  up,  we  find  that  all  the  suppositions  the  chemist 
makes  are:  That  an  atom  has  a  specific  mass,  which  is  not  altered  in 
ordinary  chemical  change,  and  consists  of  a  specific  kind  of  material; 
that  its  capacity  for  combining  with  other  atoms  is  limited  by  its 
valence,  and,  sometimes,  that  the  atoms  in  a  molecule  are  arranged 
with  reference  to  one  another  in  space  in  some  definite  way.  After 
making  this  formulative  hypothesis,  however,  he  knows  no  more 
than  he  did  before  about  the  real  mechanism  of  chemical  change. 

Exercises.  —  1.  In  previous  chapters  our  definitions  have  been 
experimental;  that  is,  have  been  expressed  in  terms  of  facts  learned 
by  experiment.  In  imitation  of  the  definitions  of  the  law  of  definite 
proportions  and  of  valence  outlined  on  p.  154,  write  out  theoretical 
definitions  of  the  following,  in  terms  of  the  atomic  hypothesis :  Phy- 
sical and  chemical  phenomenon,  multiple  proportions,  chemical  unit 
weight,  molecular  weight,  element,  compound,  symbol,  formula, 
equation. 

2.  Criticize  the  definitions:  The  atomic  weight  of  an  element  is  the 
smallest  portion  of  that  element  which  takes  part  in  chemical  change. 
An  atom  is  the  smallest  particle  that  can  be  conceived. 

3.  Define  all  the  varieties  of  chemical  change  (p.  124)  in  terms  of 
the  atomic  hypothesis. 


CHAPTER  XIV 
THE   HALOGEN   FAMILY 

THE  elements  to  which  we  have  so  far  devoted  most  attention  have 
been  oxygen,  hydrogen,  and  chlorine.  If  we  recall  the  chemical 
properties  and  relations  of  these  elements  we  shall  recognize  the  fact 
that  they  all  possess  very  distinct  individualities. 

The  Chemical  Relations  of  Elements.  —  Hydrogen  is  the  sub- 
stance (p.  74)  which  unites  readily  with  oxygen  and  chlorine,  less 
readily  with  other  non-metals,  and  scarcely  at  all  with  metals. 
Oxygen  and  chlorine  resemble  one  another  somewhat  in  the  greatness 
of  their  chemical  activity  and  the  variety  of  free  elements  with  which 
they  are  capable  of  uniting,  but  differ  markedly  in  what  we  have 
called  their  chemical  relations  (p.  116).  The  resulting  compounds 
belong,  in  fact,  to  quite  different  classes  —  oxygen  forms  oxides, 
chlorine  forms  chlorides  —  and  elements  are  considered  similar  only 
when  they  resemble  one  another  in  chemical  relations,  and  produce,  by 
combination  with  the  same  element,  compounds  having  similar  chemical 
properties.  Thus  the  common  oxide  of  hydrogen,  water,  is  a  neutral 
substance,  and  is  chemically  rather  indifferent.  The  chloride  of 
hydrogen  in  aqueous  solution  is  a  strong  acid  and  is  chemically  very 
active.  *  If  all  the  other  chemical  elements  differed  from  one  another 
as  much  as  do  these  three,  they  would  be  incapable  of  classification. 
In  reality,  however,  we  find  that  the  elements  can  be  grouped 
together  in  sets.  They  are  classified  according  to  the  kind  of  sub- 
stances with  which  they  combine  and  the  chemical  nature  of  the 
products.  In  some  families  the  resemblance  is  close,  in  others  less 
close.  The  present  group  is  of  the  former  class,  and  will  serve, 

*  The  difference  between  oxides  and  chlorides  is  seen  in  their  behavior,  and 
will  show  itself  as  we  proceed.  We  have  already  learned,  however,  that  oxides 
often  unite  with  water  to  form  acids  or  bases  (p.  81).  Chlorides  do  not 
unite  with  water  to  form  new  substances  with  marked  characteristics  (c/. 
p.  83).  They  belong  to  the  large  class  of  compounds  designated  salts  (q.v.). 

157 


158  COLLEGE   CHEMISTRY 

therefore,  as  a  convenient  beginning  in  the  work  of  tracing  relations 
between  the  elements  and  in  classifying  the  facts  of  descriptive 
chemistry. 

The  Chemical  Relations  of  the  Halogens.  —  The  bromide 
(NaBr),  iodide  (Nal),  and,  to  a  less  extent,  the  fluoride  (NaF),  of 
sodium,  resemble  sodium  chloride  (NaCl)  in  appearance  and  behavior. 
From  the  fact  that  chlorine,  bromine,  iodine,  and  fluorine  are  thus  all 
able  to  produce  substances  which,  both  in  composition  and  chemical 
behavior  (see  p.  187),  resemble  salt,  they  are  known  as  the  halogens 
(Gk.  aXs,  salt;  yewav,  to  produce),  and  their  compounds  are  named 
the  halides.  The  halogens,  as  the  above  formulae  show,  are  univa- 
lent.  They  all  form  compounds  with  hydrogen,  and  these  com- 
pounds closely  resemble  hydrogen  chloride  (q.v.).  For  example, 
they  are  colorless,  they  are  gases  (except  hydrogen  fluoride,  a  very 
volatile  liquid),  they  are  very  soluble  in  water,  and  their  solutions 
are  acids.  Other  relations  will  be  given  in  a  summary  at  the  end  of 
the  chapter. 

BROMINE  Br2. 

Occurrence.  —  The  compounds  of  chlorine,  bromine,  and  iodine 
usually  occur  together  in  nature,  while  the  compounds  of  fluorine  are 
not  found  in  the  same  sources.  Bromine  occurs  chiefly  in  the  form 
of  the  bromides  of  sodium  and  magnesium,  in  the  upper  layers  of 
the  natural  beds  of  rock  salt. 

Preparation.  —  In  the  chemical  point  of  view  there  are  three  dis- 
tinct ways  in  which  bromine  is  made.  1.  The  first  of  these  is  closely 
related  to  the  common  method  of  preparing  chlorine  (p.  112).  As 
hydrobromic  acid,  unlike  hydrochloric  acid,  is  not  formed  exten- 
sively in  connection  with  any  chemical  industry,  potassium  bromide, 
the  most  accessible  compound  of  bromine,  is  treated  with  concen- 
trated sulphuric  acid,  and  the  product  is  oxidized  with  powdered 
manganese  dioxide  in  one  operation.  The  whole  action  may  be 
represented  in  a  single  equation  (see  next  section),  as  follows: 

2KBr  +  3H2SO4  +  Mn02  ->  MnSO4  +  2KHS04  +  2H2O  +  Br2. 

Bromine  being  a  volatile  liquid,  while  the  two  sulphates  are  invol- 
atile,  its  vapor  passes  off  along  with  a  little  water  when  the  above 
mixture  is  heated.  It  is  condensed  in  a  worm-tube  surrounded  by 
cold  water. 


BROMINE  !.-«) 

2.  The  second  method  of  preparing  bromine  depends  on  the  fact 
that  chlorine  is  a  more  active  element  and  displaces  bromine  from 
combination.     When,  therefore,  chlorine  is  passed  into  a  solution  of 
potassium  or  sodium  bromide,  potassium  or  sodium  chloride  is  formed 
and  the  bromine  liberated : 

2NaBr  +  C12  ->  2NaCl  +  Br2. 

When  the  liquid  is  warmed,  the  bromine  passes  off  along  with  a  part 
of  the  water,  and  may  be  condensed  as  before. 

3.  Aqueous  solutions  of  soluble  bromides  may  be  decomposed  by 
means  of  a  current  of  electricity.     The  bromine  is  set  free  at  the 
positive  electrode. 

The  whole  of  the  bromine  used  in  commerce  is  at  present  manu- 
factured in  the  first  two  of  these  ways.  Two-thirds  of  the  supply  is 
obtained  from  Stassfurt,  where,  after  the  extraction  of  the  potassium 
chloride  from  the  impure  carnallite  (KC1,  MgCl2,  6H20),  the  mother- 
liquor  is  found  to  contain  the  more  soluble  sodium  and  magnesium 
bromides  in  considerable  quantities.  The  warm  mother-liquor 
trickles  down  over  round  stones  in  a  tower.  The  chlorine  is  intro- 
duced from  below  and  dissolves  in  the  liquid.  The  bromine  is  thus 
liberated  and  passes  off  as  vapor.  A  part  of  our  supply  of  bromine 
is  obtained  from  the  brines  of  Ohio,  West  Virginia,  and  Kentucky. 
Here  the  liquid,  after  most  of  the  common  salt  has  been  removed  by 
crystallization,  is  assayed  to  ascertain  the  quantity  of  bromine  which 
it  contains,  and  is  treated  with  the  calculated  amount  of  sulphuric 
acid  necessary  for  the  action.  Manganese  dioxide  is  then  added  in 
small  quantities  at  a  time.  In  Michigan  the  brines  are  treated  with 
electrolytic  chlorine.  The  quantity  produced  in  America  in  1906, 
was  640  tons,  and  was  valued  at  $165,204. 

Partial  Equations,  a  Plan  for  Making  Complex  Equations. 

—  When  an  equation  involves  more  than  two  initial  substances  or 
products,  as  does  the  one  given  above  for  the  first  method  of  pre- 
paring bromine,  it  cannot  readily  be  worked  out  by  the  method 
formerly  recommended  (p.  41).  After  the  formulae  of  all  the  sub- 
stances, on  both  sides,  have  been  set  down,  it  is  difficult  to  hit  upon 
the  proper  coefficients  required  to  balance  the  equation.  In  such 
cases,  a  good  plan  is  to  select  two  of  the  initial  substances,  and  make 
a  partial  equation  showing  part  of  the  action  and  including  at  least 
one  actual  product.  Any  unused  units  (not  constituting  a  product) 


160  COLLEG$   CHEMISTRY 

are  then  set  down  also  and  treated  as  a  balance.  Thus  the  first  two 
of  the  substances,  in  the  above  equation,  will  furnish  the  second  of 
the  products: 

KBr  +  H2S04  ->  KHS04  (+  HBr).  (1) 

Then  we  may  represent  the  formation  of  the  first  of  the  products 
from  the  materials  which  evidently  must  have  been  used  up  in  form- 
ing it,  and  set  down  the  balance  as  before: 

MnO2  +  H2SO4  -»MnSO4  +  H2O  (+  O).  (2) 

We  then  perceive  that  the  last  of  the  products  might  come  from  the 
oxidation  of  the  first  balance  by  the  second: 

(2HBr)  4-  (O)  ->  H2O  4-  Br2.  (3) 

The  third  partial  equation  shows  that  2HBr  will  be  needed  for  the 
amount  of  O  obtainable  from  MnO2,  so  we  go  back  to  (1)  and  multiply 
it  by  two  throughout: 

2KBr  +  2H2S04  -+  2KHSO4  (+  2HBr)  (1) 

MnO2  +  H2SO4  ->  MnS04  +  H2O  (+  0)  (2) 

(2HBr)  +  (0)  ->  H2O  +  Br2 (3) 

2KBr  +  3H2SO4  +  MnO2  ->  2KHSO4  +  MnSO4  +  2H2O  +  Br2 

When  we  now  add  the  real  substances  used  and  produced,  as  they 
occur  in  these  partial  equations,  and  leave  out  the  balances,  which 
have  been  adjusted  so  as  to  cancel  one  another,  we  obtain  the  final 
equation  already  given  for  the  action.  It  must  be  observed  that  this 
subdivision  of  the  action  into  parts  is  a  purely  arithmetical  device. 
The  partial  equations  made  for  purposes  like  the  present  one  are 
often  purely  fictitious.  It  is  still  true,  however,  that  we  are  aided 
in  the  selection  of  partial  actions  at  each  step  by  following  some 
plausible  theory  as  to  stages  for  the  action  which,  if  there  were  any, 
would  be  chemically  conceivable. 

Physical  Properties.  —  Bromine  is  a  dark-red  liquid  (sp.  gr. 
3.18).  It  boils  at  59°,  forming  a  deep-red  vapor,  and  even  at  ordi- 
nary temperatures  gives  a  high  vapor  pressure  (150  mm.  at  18°)  and 
evaporates  quickly.  When  cooled  it  forms  red,  needle-shaped  crys- 
tals (m.-p.  — 7.3°).  A  saturated  aqueous  solution  (bromine-water) 
at  ordinary  temperatures  contains  about  three  parts  of  bromine  in 
one  hundred  of  water.  The  element  is  much  more  soluble  in  carbon 


HYDROGEN   BKOMIDi;  1(51 

> 

bisulphide,  alcohol,  and  other  organic  solvents.  Its  vapor  density 
up  to  750°  is  160  (oxygen  =  32). 

Bromine  (Gk.  /fyw/xos,  a  stench)  has  a  most  pungent  odor.  It  has 
a  very  irritating  effect  on  the  mucous  membrane  of  the  nostrils  and 
throat.  If  spilled  upon  the  hands  it  has  a  most  destructive  action 
upon  the  tissues. 

When  free  bromine  is  added  to  starch  emulsion  no  special  change 
in  tint*  is  observable  unless  a  very  large  amount  of  the  element  is 
used  (see  Iodine). 

Chemical  Properties.  —  The  molecular  weight  of  bromine  being 
160,  and  the  atomic  weight  79.96,  the  molecules  of  the  vapor  are 
diatomic  and  the  molecular  formula  of  the  elementary  substance 
is  Br2. 

Bromine  unites  directly  with  hydrogen.  The  mixture  of  the  gases 
is  not  explosive,  and  the  union  is  much  slower  than  in  the  case  of 
chlorine.  In  presence  of  finely  divided  platinum  the  speed  of  the 
action  may  be  considerably  increased. 

Bromine  forms  compounds  directly,  both  with  non-metals,  like 
phosphorus  and  arsenic,  and  with  most  of  the  metals.  Towards 
unsaturated  substances  and  organic  compounds  it  behaves  like 
chlorine  (q.v.).  In  all  cases  the  interaction  is  less  violent  than  when 
chlorine  is  used,  and  the  element  is  displaced  from  combination  with 
hydrogen  and  with  the  metals  by  free  chlorine. 

Potassium  bromide  is  employed  in  making  photographic  plates 
and  in  medicine  and  in  the  preparation  of  organic  dyes. 

HYDROGEN  BROMIDE  HBr. 

Preparation.  —  It  might  be  expected  that  the  most  convenient 
way  of  producing  this  compound  would  be  similar  to  that  used  in 
preparing  hydrogen  chloride,  namely,  by  the  action  of  concentrated 
sulphuric  acid  upon  some  common  bromide,  such  as  potassium  bro- 
mide (KBr  +  H2SO4  <±  HBr  +  KHSO4).  Hydrogen  bromide  being 
less  stable,  however,  a  large  part  of  it  is  oxidized  by  the  sulphuric 
acid  and  the  product  is  mixed  with  sulphur  dioxide  and  free  bromine. 

Since  all  acids  decompose  all  salts  more  or  less,  use  of  an  acid  which 
does  not  give  up  its  oxygen  so  readily,  such  as  phosphoric  acid,  will 
yield  pure  hydrogen  bromide  (KBr  +  H8PO4->HBr|  +  KH2P04). 
The  small  solubility  of  the  salt  in  concentrated  phosphoric  acid 


162 


COLLEGE    CHEMISTRY 


retards  the  interaction  and  makes  the  evolution  of  the  gas  very  slow, 
however. 

Pure  hydrogen  bromide  is  best  prepared  by  the  action  of  water  upon 
phosphorus  tribromide  (see  Hydrolysis,  below) .  When  bromine  and 
phosphorus  are  mixed,  a  violent  union  of  the  two  elements  takes 
place,  producing  phosphorus  tribromide  (PBr3).  This  substance, 
which  is  a  colorless  liquid,  is  in  turn  broken  up  with  great  ease  by 
water,  producing  phosphorous  acid,  which  is  not  volatile,  and 
hydrogen  bromide: 

,  Br       HOH 

P— Br  +  HOH  ->  P(OH)3  +  3HBr. 
NBr       HOH 

In  practice,  those  two  actions  are  carried  on  simultaneously.     To 
diminish  the  vigor  of  the  interaction,  red  phosphorus  is  taken  instead 


FIG.  47. 

of  yellow,  and  is  mixed  with  two  or  three  times  its  weight  of  sand  in 
a  flask  (Fig.  47).  A  small  quantity  of  water  is  added.  Excess  of 
water  must  be  avoided,  as  the  hydrogen  bromide  produced  is 
extremely  soluble,  and  would  therefore  be  retained  in  the  flask 
instead  of  being  disengaged  as  gas.  The  bromine  is  placed  in  the 
dropping  funnel,  and  admitted,  a  little  at  a  time,  to  the  mixture. 
The  gas  produced  is  passed  through  a  U-tube  containing  glass  beads 
mixed  with  red  phosphorus.  The  latter  combines  with  any  free 
bromine  which  may  have  been  carried  along  with  the  gas.  The 


HYDROGEN    BROMIDE  163 

second  U-tube,  containing  water,  may  be  attached  when  a  solution 
of  the  gas  is  required. 

Hydrolysis.  —  The  interaction  of  water  with  phosphorus  tri- 
bromide  (foregoing  section)  illustrates  an  important  property  of 
water  (p.  81)  and  a  common  variety  of  chemical  change.  The 
action  is  a  double  decomposition  in  which  water  is  one  of  the  interacting 
substances  and  is  called  an  hydrolysis  (Gk.  vSup,  water,  and  Awris,  the 
act  of  loosing).  The  water  divides  into  the  radicals  H  and  OH,  and 
the  former  unites  with  the  more  active  non-metal  in  the  substance 
(as  the  bromine  in  PBr3)  and  the  hydroxyl  with  the  other  element. 
For  example,  PC13  +  3HOH  -»  P(OH)3  +  3HC1.  All  the  halides 
of  the  non-metals  are  thus  hydrolyzed  (see  sulphur  mono-chloride), 
as  are  also  some  other  classes  of  compounds. 

Physical  Properties.  —  Hydrogen  bromide  is  a  colorless  gas 
with  a  sharp  odor.  It  is  two  and  a  half  times  as  heavy  as  air.  It  is 
easily  reduced  to  the  liquid  condition.  It  is  exceedingly  soluble  in 
water,  and  in  contact  with  moist  air  condenses  the  water  vapor  to 
clouds  of  liquid  particles.  Pure  hydrogen  bromide,  whether  in  the 
gaseous  condition  or  in  the  liquefied  form,  is  a  nonconductor  of 
electricity  (see  below). 

Chemical  Properties.  —  The  chemical  properties  of  hydrogen 
bromide  are  similar  to  those  of  hydrogen  chloride  (p.  120).  It  is 
somewhat  less  stable,  and  dissociation  into  its  constituents  begins  to 
be  noticeable  at  800°.  When  free  from  water,  it  is  not  an  acid  (see 
below) .  The  gas  interacts  vigorously  with  chlorine,  hydrogen  chloride 
and  free  bromine  being  produced,  2HBr  +  C12  — >  2HC1  +  Br2. 

Chemical  Properties  of  Hydrobromic  Acid  HBr,  Aq.  —  The 

solution  of  the  hydrogen  bromide  in  water  is  an  active  acid  (cf. 
p.  75).  It  conducts  electricity  extremely  well.  In  contact  with 
certain  metals,  and  with  oxides  of  metals  and  hydroxides  of  metals, 
it  behaves  exactly  like  hydrochloric  acid  (p.  123).  In  the  first  case, 
hydrogen  is  set  free  and  the  bromide  of  the  metal  produced.  In  the 
other  two  cases,  water  and  the  bromides  of  the  metals  are  produced. 
Oxidizing  agents  set  bromine  free  from  hydrobromic  acid,  and  the 
only  difference  as  compared  with  hydrochloric  acid  is  that  less 
powerful  oxidizing  agents  can  produce  this  result  (p.  161).  Chlorine 
dissolved  in  water  displaces  bromine  from  hydrobromic  acid  and 
from  soluble  bromides  with  ease. 


164  COLLEGE    CHEMISTRY 

IODINE  I2. 

Occurrence.  —  Iodine,  like  bromine,  occurs  in  sea-water,  although 
it  is  present  in  much  smaller  quantities.  Fortunately,  certain  species 
of  sea-weed  seem  to  have  the  power  to  remove  it  from  water,  and 
use  it  as  a  constituent  in  complex  organic  compounds  which  they 
contain.  In  the  case  of  some  species,  the  ash  of  the  sea-weed  is  said 
to  contain  as  much  as  two  per  cent,  or  even  more.  The  other  chief 
source  of  iodine  is  in  Chili  saltpeter  (mainly  NaNO3),  in  which  it  is 
present  in  the  form  of  small  proportions  of  sodium  iodate  (NaIO3) 
and  sodium  iodide.  Of  recent  years  the  quantity  obtained  commer- 
cially from  this  source  has  greatly  increased,  while  that  from  sea- 
weed has  diminished. 

Preparation. —  l.  In  factories  where  the  iodine  is  extracted 
from  sea-weed,  the  latter  is  first  burned,  either  roughly  in  hollows 
in  the  ground,  or  more  carefully  in  specially  constructed  ovens.  A 
great  deal  of  the  iodine  seems  to  be  lost  by  volatilization  in  this  pro- 
cess, but  the  ash  that  remains,  which  is  known  in  Scotland  as  kelp 
and  in  Normandy  as  varec,  still  contains  from  0.5  to  1.5  per  cent  of 
sodium  iodide.  The  ash  is  treated  with  water,  and  the  solution  is 
evaporated  so  as  to  permit  the  deposition  of  the  sodium  chloride  and 
sodium  sulphate  which  it  contains.  The  sodium  iodide,  being  very 
soluble,  remains  in  the  mother-liquor.  This  is  then  treated  with 
manganese  dioxide  and  sulphuric  acid.  The  quantity  of  manganese 
dioxide  is  carefully  measured  so  as  to  be  just  sufficient  to  set  free 
the  iodine  contained  in  the  liquid,  without  proceeding  farther  to  the 
liberation  of  the  chlorine  which  it  contains  in  much  larger  amounts. 
When  the  mixture  is  heated,  the  iodine  passes  off  in  the  form  of 
vapor,  and  is  condensed  in  a  suitable  receiver.  The  action  (cf.  pp. 
113,  158,  160)  is: 

2NaI  +  MnO2  +  3H2SO4  ->  MnSO4  +  2NaHSO4  +  2H20  -{-  I2. 

2.  In  France  the  treatment  is  similar,  excepting  that  chlorine  is 
used  to  liberate  the  iodine  in  the  last  stage  (2NaI  +  Cl2^2NaCl  +  I2). 
The  quantity  is  adjusted  so  that  excess  may  not  be  employed.     The 
iodine,  being  insoluble,  forms  a  dense  precipitate,  and,  when  the 
liquid  is  pressed  out,  it  remains  behind  in  the  form  of  a  paste. 

3.  Electricity  could  also  be  used  for  the  decomposition  of  this 
mother-liquor.    The  iodine  is  set  free  at  the  positive  electrode,  while 
the  metals  pass  to  the  other. 


IODINE  165 

In  all  cases  the  iodine  is  subjected  to  purification  before  being 
sold.  This  is  carried  out  by  distilling  it  with  a  little  powdered 
potassium  iodide.  It  condenses  in  the  solid  form  directly,  in  glitter- 
ing black  plates  (sublimed  iodine).  The  distillation  of  a  solid  body, 
when  a  condensation  takes  place  directly  to  the  solid  form,  is  spoken 
of  as  sublimation. 


Physical  Properties.—  Iodine  (Gk.  toaSifc,  like  a  violet)  is  a 
solid  substance  (sp.  gr.  5),  exhibiting  large  crystalline  plates  of 
rhombic  form.  It  melts  at  114°,  and  boils  at  184°.  The  vapor  has 
at  first  a  reddish-violet  tint,  and  on  being  more  strongly  heated 
becomes  deep  blue  (see  next  section)  . 

Iodine  is  much  less  soluble  in  water  than  are  the  other  halogens, 
and  the  solution  has  a  scarcely  perceptible  brown  tint.  At  ordinary 
temperatures  one  part  of  iodine  dissolves  in  5000  —  6000  parts  of 
water.  It  is  much  more  soluble  in  carbon  disulphide  (p.  103)  and 
in  chloroform,  in  which  it  gives  violet  solutions.  In  alcohol  it  gives 
a  solution  which  is  brown.  The  brown  color  is  attributed  to  the 
fact  that  in  alcohol  the  iodine  is  in  a  condition  of  feeble  combination, 
and  not  simply  in  solution.  An  aqueous  solution  of  potassium  iodide, 
hydrogen  iodide,  or  any  other  iodide,  has  likewise  the  power  to  take 
up  large  quantities  of  iodine.  In  these  cases,  however,  the  formation 
of  definite  compounds  (such  as,  KI  +  I2  +±  KI3),  by  a  reversible 
action,  accounts  for  the  amount  of  iodine  which  appears  to  be 
dissolved. 

The  behavior  of  free  iodine  towards  starch  forms  a  distinctive  test 
for  both  substances  (cf.  p.  114).  When  starch  is  treated  with  boiling 
water  it  passes  into  a  state  of  fine  suspension,  and  the  liquid  may  be 
filtered  without  removal  of  all  the  starch.  When  traces  of  iodine, 
contained,  for  example,  in  the  pale-brown  aqueous  solution,  are 
added  to  this  filtered  starch  emulsion,  a  deep-blue  color  is  produced. 
The  same  action  is  used  as  a  test  for  starch.  This  blue  substance  is 
not  a  chemical  compound,  but  a  solution  of  the  iodine  in  the  solid 
starch  which  is  suspended  in  the  water. 

Chemical  Properties.  —  The  molecular  weight  of  iodine,  ascer- 
tained by  weighing  the  vapor  at  a  temperature  above  the  boiling- 
point,  is  254.8.  The  atomic  weight  being  126.97,  the  molecule 
contains  two  atoms.  This  value  for  the  density.  remains  unchanged 
when  the  measurement  is  made  at  temperatures  up  to  about  700°. 


166  COLLEGE   CHEMISTRY 

Beyond  this  point,  however,  the  vapor  diminishes  in  density  more 
rapidly  than  Charles'  law  would  lead  us  to  expect,  and  at  1700°  the 
molecular  weight  has  fallen  to  127.  As  the  vapor  is  heated,  a  larger 
and  larger  proportion  of  the  molecules  is  broken  up,  until  the  decom- 
position has  become  complete.  As  in  all  cases  of  dissociation,  when 
the  vapor  is  cooled  the  atoms  recombine  to  form  molecules.  This 
chemical  action  is  interesting,  since  it  is  the  most  notable  case  in 
which  we  encounter  both  the  monatomic  and  the  diatomic  forms  of 
the  same  element.  The  heat  given  out  when  the  atoms  reunite  to 
form  the  molecules  is  very  considerable  (21  <=»  I2  +  28,500  cal.),  indi- 
cating that  the  chemical  union  of  two  atoms  of  identical  nature  may 
be  as  vigorous  as  that  of  two  atoms  of  different  chemical  substances. 
The  monatomic  and  diatomic  forms  of  iodine  should  be  distinct 
chemical  substances,  and  if  the  investigation  of  the  behavior  of  the 
former  were  not  hampered  by  the  very  high  temperature  at  which 
alone  it  exists,  it  would  doubtless  be  found  to  exhibit  different  chem- 
ical properties. 

Iodine  forms  no  hydrate  with  water.  It  unites  very  slowly  with 
hydrogen,  even  when  heated.  Iodine  unites  directly  with  a  number 
of  elements,  including  some  non-metals  and  the  majority  of  the 
metals.  When  phosphorus  is  presented  in  the  yellow  form,  the  action 
takes  place  spontaneously  without  the  assistance  of  heat.  Both 
chlorine  and  bromine  displace  iodine  from  combination  with  hydro- 
gen and  the  metals  (2HI  4-  Br2  ->  2HBr  +  I2).  The  action  may  be 
brought  about  either  with  the  substances  in  dry  form  or  with  their 
aqueous  solutions. 

Iodine  and  its  compounds  are  much  used  in  the  arts  and  medicine. 
Iodine  is  applied,  in  the  form  of  an  alcoholic  solution  ("  tincture  of 
iodine  ");  for  the  reduction  of  some  swellings.  It  is  required  in  mak- 
ing iodoform  (CHI3),  and  the  iodides  of  potassium,  rubidium,  and 
sodium,  which  are  used  in  medicine.  Potassium  iodide  is  likewise 
employed  in  the  manufacture  of  photographic  plates. 

HYDROGEN  IODIDE  HI. 

Preparation.  —  The  direct  union  of  hydrogen  and  iodine  cannot 
be  employed  in  preparing  pure  hydrogen  iodide.  The  union  takes 
place  slowly  at  445°,  but,  since  the  action  is  markedly  reversible 
(c/.p.  115),  always  remains  incomplete:  H2  +  I2?=±2HI  (see  p.  173). 

The  action  of  concentrated  sulphuric  acid  upon  potassium  or 


HYDROGEN   IODIDE  167 

sodium  iodide  is  equally  inapplicable.  In  this  case,  as  in  that  of 
hydrogen  bromide  (p.  161),  the  hydrogen  halide  reduces  the  sul- 
phuric acid,  and  much  free  iodine  is  formed. 

Finely  powdered  sodium  iodide  and  concentrated  rJhosphoric  acid 
(cf.  p.  161),  when  mixed  and  warmed,  give  pure  hydrogen  iodide 
(Nal  +  H3P04^HIt  +  NaH2P04).  This  action,  although  it 
goes  very  slowly,  was  formerly  used  in  preparing  the  gas. 

The  best  method  is  one  similar  to  that  described  under  hydrogen 
bromide.  Phosphorus  and  iodine  unite  directly  to  form  PI3.  This 
is  a  yellow  solid  which  is  violently  decomposed  by  water  and  gives 
phosphorous  acid  and  hydrogen  iodide: 

PI3  +  3H20  ->  P(OH)3  +  3HL 

If  excess  of  water,  which  dissolves  hydrogen  iodide,  is  avoided,  the 
latter  goes  off  in  a  continuous  stream  in  a  gaseous  condition. 

Still  another  method  of  making  hydrogen  iodide  is  frequently 
employed  when  a  solution  of  the  gas  in  water  is  required,  and  not  the 
gas  itself.  Powdered  iodine  is  suspended  in  water,  and  hydrogen 
sulphide  gas  (q.v.)  is  introduced  through  a  tube  in  a  continuous 
stream.  The  iodine  dissolves  slowly  in  the  water,  I2  (solid)  +±  I2 
(dslvd),  and  acts  upon  the  hydrogen  sulphide,  which  likewise  dis- 
solves, H2S  (gas)  <=±  H2S  (dslvd).  Sulphur  separates  in  a  fine  powder, 
S  (dslvd)  <=±  S  (solid),  and  hydrogen  iodide  is  formed  in  accordance 
with  the  equation: 

H2S  +  I2  ->  2HI  +  S. 

This  action  takes  place,  however,  only  in  presence  of  water,  although 
the  water  does  not  appear  in  the  equation.  The  solution  is  freed 
from  the  deposit  of  sulphur  by  filtration,  and  may  be  concentrated  to 
57  per  cent  of  hydriodic  acid  by  distilling  off  the  water. 

Physical  Properties.  —  Hydrogen  iodide  is  a  colorless  gas  with 
a  sharp  odor.  Its  molecular  weight  is  128,  and  it  is  therefore  much 
heavier  than  air,  the  average  weight  of  whose  molecules  is  28.955 
(p.  126).  It  is  a  nonconductor  of  electricity,  both  in  the  g 
and  in  the  liquefied  conditions.  It  is  exceedingly  soluble  in 
so  that  at  0°  ten  grams  of  water  will  absorb  ninety  grams  of  the  gas, 
giving  a  90  per  cent  solution.  The  behavior  of  this  solution  is  simi- 
lar to  those  of  hydrogen  chloride  and  hydrogen  bromide  (cf.  p.  120). 
The  mixture  of  constant  boiling-point  distils  over  at  127°  (at  760 
mm.),  and  contains  57  per  cent  of  hydrogen  iodide. 


168  COLLEGE   CHEMISTRY 

Chemical  Properties. —  Hydrogen  iodide  is  the  least  stable  of  the 
hydrogen  halides.  When  heated  it  begins  visibly  to  decompose  into 
its  constituents  at  180°.  On  account  of  the  ease  with  which  it  parts 
with  the  hydrogen  which  it  contains,  it  can  be  burned  in  oxygen 
gas,  4HI  +  O2  — »  2H20  +  2I2.  When  the  gas  is  mixed  with  chlorine, 
a  violent  chemical  change,  accompanied  by  a  flash  of  light,  occurs, 
the  iodine  is  set  free,  and  hydrogen  chloride  is  produced,  C12  +  2HI— » 
2HC1  +  I2.  Bromine  vapor  will  similarly  displace  the  iodine  from 
hydrogen  iodide. 

Chemical  Properties  of  Hydriodic  Acid  HI,  Aq.  —  In  most 
respects  the  aqueous  solution  behaves  exactly  like  hydrochloric  and 
hydrobromic  acids.  With  oxidizing  agents,  for  example,  such  as 
manganese  dioxide,  it  gives  free  iodine,  just  as  the  others  give  free 
chlorine  and  bromine,  respectively.  Here,  however,  the  oxidation  is 
so  much  more  easily  carried  out,  that  it  is  slowly  effected  by  atmos- 
pheric oxygen,  so  that  hydriodic  acid  left  exposed  to  the  air  gradually 
becomes  brown  (02  +  4HI  ~>  2H20  +  2I2). 

Although  the  dry  gas  is  not  an  acid,  the  solution  has  all  the  ordi- 
nary properties  of  this  class  of  substances  (cf.  pp.  63,  111,  123,  161). 
The  hydrogen  may  be  displaced  by  metals  like  zinc  and  magnesium. 
The  acid  interacts  wiih  oxides  and  hydroxides,  forming  iodides  and 
water. 

FLUORINE  F2. 

The  discussion  of  this  element  should  logically  have  preceded  that 
of  chlorine,  since  it  is,  of  all  the  members  of  the  halogen  family,  the 
most  active.  Chlorine  was  taken  up  first,  however,  because  its  com- 
pounds are  more  familiar.  Fluorine  is  found  in  combination  in 
nature.  It  occurs  chiefly  in  the  mineral  fluorite  (calcium  fluoride, 
CaF2)  and  in  cryolite,  a  double  fluoride  of  aluminium  and  sodium 
(3NaF;  A1F3). 

Preparation.  —  When  a  solution  of  hydrofluoric  acid  is  heated 
with  manganese  dioxide,  oxidation  does  not  occur  and  free  fluorine 
is  not  produced.  Until  very  recently  all  efforts  to  isolate  the  element 
failed.  It  was  perfectly  understood  that  the  reason  of  these  failures 
lay  in  the  greater  chemical  activity  of  fluorine,  which  made  it  more 
difficult  of  separation  from  any  state  of  combination  than  the  other 
halogens.  Its  preparation  was  finally  achieved  by  Moissan  (1886) 


FLUORINE 


by  the  decomposition  of  anhydrous  hydrogen  fluoride,  which  is 

liquid  below  19°,  by  means  of  electricity.     The  apparatus  (Fig.  48) 

is  made  of  copper,  which,  after  receiving  a  thin  coating  of  the  fluoride, 

is  not   further   affected.     To  reduce  the 

tendency  to  chemical  union,  the  whole  is 

immersed  in  a  bath  giving  a  temperature 

of  —  23°.     The  electrodes  are  made  of  an 

alloy  of  platinum  and  iridium,  which  is  the 

only  substance  that  can  resist  the  action 

of  the  fluorine.     Hydrogen  fluoride,  like 

other  hydrogen  halides,  is  a  nonconductor 

of  electricity,    and   a   small   quantity   of 

potassium  fluoride   has   to   be   added  to 

enable  the  current  of  electricity  to  pass. 

The    fluorine  is  set  free  at  the  positive 

electrode,    and  hydrogen  appears  at  the 

negative.     The  U-tube  is  closed,  after  the 

introduction  of  the  hydrogen  fluoride,  by 

means  of  blocks  made  of  calcium  fluoride, 

which  is  naturally  unable  further  to  enter 

into  combination  with  fluorine.     For  the 

reception  and  examination  of  the  fluorine 

gas,  other  copper  tubes  can  be  screwed  on  to  the  side  necks  of  the 

apparatus,  and,  when  necessary,  small  windows  of  calcium  fluoride 

can  be  provided.      It  has  been  found  that  fluorine  dried  with 

extraordinary  precautions  is  without  action  on  glass. 

Physical  Properties.  —  Fluorine  is  a  gas  whose  color  is  like  that 
of  chlorine,  but  somewhat  paler.  Its  density  (38)  shows  that  the 
molecule  is  diatomic  (F2).  The  gas  is  the  most  difficult  of  the  halogens 
to  liquefy.  The  liquid  boils  at  -  186°. 

Chemical  Properties.  • —  Fluorine  unites  with  every  element, 
with  the  exception  of  oxygen  and  the  members  of  the  helium  family, 
and  in  many  cases  does  so  with  such  vigor  that  the  union  begins 
spontaneously  without  the  assistance  of  external  heat.  Dry  plati- 
num and  gold  are  the  elements  least  affected.  It  explodes  with 
hydrogen  at  the  ordinary  temperature,  without  the  assistance  of 
sunlight.  Fluorine  displaces  oxygen  from  water  instantaneously 


FIG.  48. 


170  COLLEGE   CHEMISTRY 

and  gives  ozone  (q.v.}.  On  the  introduction  of  a  drop  of  water 
into  a  tube  of  fluorine,  the  vessel  is  filled  with  the  deep-blue  gas, 
3F2  +  3H20  ->  3H2F2  +  03. 

Fluorine  displaces  the  chlorine  in  hydrogen  chloride  as  easily  as 
chlorine  in  turn  displaces  bromine  or  iodine. 

HYDROGEN  FLUORIDE  H2F2. 

Preparation.  —  Pure,  dry  hydrogen  fluoride  is  best  made  by 
heating  potassium-hydrogen  fluoride,  2KHF2  <=±  K2F2  +  H2F2  f . 
For  ordinary  purposes,  however,  the  preparation  of  an  aqueous 
solution  is  the  ultimate  object.  Usually  powdered  calcium  fluoride 
is  treated  with  concentrated  sulphuric  acid,  and  the  mixture  distilled 
in  a  platinum  retort: 

CaF2  +  H2SO4<=»CaS04  +  H2F2|. 

The  hydrofluoric  acid  passes  over  and  is  caught  in  distilled  water. 
The  aqueous  solution  thus  obtained  has  to  be  kept  in  vessels  made 
of  lead,  rubber,  or  paraffin,  as  glass  interacts  with  the  acid  with  great 
rapidity  (see  below). 

Physical  Properties.  —  Hydrogen  fluoride  is  a  colorless  liquid, 
boiling  at  19.4°.  It  mixes  freely  with  water,  and,  on  distillation,  an 
acid  of  constant  boiling-point  (120°  at  760  mm.)  containing  35  per 
cent  of  hydrogen  fluoride  is  obtained.  The  vapor  density  of  the 
hydrogen  fluoride  between  its  boiling-point  (19.4°)  and  30°  corres- 
ponds to  a  molecular  weight  of  40,  and  the  formula  should  therefore 
be  H2F2.  Above  this  temperature  the  vapor  becomes  gradually 
lighter  (p.  146)  and,  when  88°  is  reached,  the  molecular  weight  has 
fallen  to  20,  corresponding  to  the  formula  HF. 

Chemical  Properties  of  Hydrofluoric  AcidH2F29  Aq.  —  Metals 
like  zinc  and  magnesium  interact  with  hydrofluoric  acid  with  evolu- 
tion of  hydrogen.  The  action  is  less  violent  than  with  other  halogen 
acids.  The  acid  interacts  with  oxides  and  hydroxides,  forming 
fluorides.  The  chief  difference  in  this  respect  which  it  exhibits, 
when  compared  with  the  other  halogen  acids,  is  one  which  we  should 
expect  from  its  formula,  H2F2.  We  may  displace  either  one  or  both 
the  hydrogen  atoms  in  the  molecule  with  a  metal.  Thus,  one  of  the 
commonest  salts  of  hydrofluoric  acid  is  potassium-hydrogen  fluoride 


HALOGENS  AS  A  FAMILY  171 

KHF2  (q.v.)  mentioned  above.  In  this  respect  the  acid  resembles 
sulphuric  acid  and  other  acids  containing  more  than  one  replaceable 
hydrogen  unit.  Salts  in  which  a  portion  of  the  acid  hydrogen  still 
remains  undisplaced  are  spoken  of  as  acid  salts. 

The  most  remarkable  property  of  hydrofluoric  acid  depends  on 
the  great  tendency  which  fluorine  has  to  unite  with  silicon,  forming 
the  gaseous  silicon  tetrafluoride.  Glass  (q.v.)  is  a  mixture  of  silicates 
of  calcium  and  sodium,  and  is  rapidly  decomposed  by  hydrofluoric 
acid.  The  nature  of  the  change  is  shown  by  the  two  following 
equations: 

CaSiO3  +  3H2F2  ->  SiF4  +  CaF2    +  3H2O, 
Na2SiO3  +  3H2F2  ->  SiF4  +  Na2F2  +  3H2O. 

All  other  silicates  are  decomposed  according  to  the  same  plan.  The 
silicon  tetrafluoride  is  a  gas.  The  fluorides  of  calcium  and  sodium 
are  solid  and  crumble  away  or  dissolve.  Thus  the  glass  is  completely 
disintegrated.  The  vapor  of  hydrofluoric  acid,  generated  in  the  way 
described  above  from  calcium  fluoride  in  a  lead  dish,  is  used  for 
etching  glass.  The  surface  of  the  glass  is  covered  with  paraffin  to 
protect  it  from  the  action  of  the  vapor,  and  with  a  sharp  instrument 
portions  of  this  paraffin  are  removed  where  the  etching  effect  is 
desired.  The  vapor  gives  a  rough  surface  where  it  encounters  the 
glass.  The  aqueous  solution,  which  may  also  be  employed,  makes 
smooth  depressions  on  the  surface. 

THE  HALOGENS  AS  A  FAMILY. 

It  may  be  useful  here  to  bring  together  some  of  the  facts  in  regard 
to  the  halogens  and  their  compounds  by  way  of  showing  more  clearly 
how  far  they  resemble  one  another,  and  in  what  ways  they  differ. 
The  most  noticeable  fact  is  that,  if  we  arrange  them  in  order  in  respect 
to  any  one  property,  chemical  or  physical,  the  other  properties  will 
be  found  to  place  them  in  the  same  order.  Thus,  if  we  consider  (1) 
the  physical  properties,  we  find  that  the  color  deepens  as  we  pass 
from  fluorine  through  chlorine  and  bromine  to  iodine.  The  specific 
gravities  of  the  elements  increase  in  the  same  order.  The  volatility 
of  the  elements  decreases  in  the  same  way  —  fluorine  being  the  hard- 
est to  liquefy,  while  iodine  is  a  solid  and  boils  at  a  fairly  high  tem- 
perature. (2)  In  their  chemical  behavior,  when,  for  example,  they 
unite  with  the  metals  and  hydrogen,  the  vigor  of  the  action  is  greatest 


172  COLLEGE   CHEMISTRY 

with  fluorine  and  diminishes  progressively  until  we  reach  iodine. 
We  shall  see  later  that  the  affinity  for  oxygen,  on  the  other  hand, 
increases  as  we  pass  from  fluorine  to  iodine. 

3.  The  relations  of  these  elements  in  combination  show  that  they 
are  all  univalent  in  respect  to  union  with  hydrogen  and  metals.  In 
their  oxygen  compounds  (q.v.),  however,  they  frequently  exhibit  a 
higher  valence.  The  compounds  which  they  form  with  any  one 
element  are  usually  very  similar  to  one  another.  All  their  com- 
pounds with  hydrogen,  for  example,  become  acids  when  dissolved  in 
water.  The  oxides  interact  with  water  to  give  acids,  and  the  halo- 
gens are  therefore  non-metals  (p.  82).  The  most  noticeable  lack 
of  harmony  in  this  group  is  observed  when  we  consider  the  solubili- 
ties of  the  corresponding  compounds.  Thus,  the  potassium  salts  are 
all  soluble  in  water.  Silver  chloride,  bromide,  and  iodide  are  almost 
insoluble,  the  amount  dissolved  decreasing  in  that  order.  Silver 
fluoride,  however,  is  quite  soluble.  Calcium  chloride,  bromide,  and 
iodide  are  all  very  soluble,  while  calcium  fluoride  is  almost  completely 
insoluble. 

It  will  be  noted  that  the  order  in  which  the  halogens  are  thus 
placed  by  consideration  of  most  of  their  properties  is  the  order  of 
increasing  atomic  weights  (see  Periodic  system). 

COMPOUNDS  OF  THE  HALOGENS  WITH  EACH  OTHER. 

Iodine  unites  directly  with  chlorine  to  form  two  compounds.  The 
more  familiar  one  is  a  red  crystalline  substance,  iodine  monochloride 
IC1.  Another  compound,  IC13,  is  made  by  the  use  of  excess  of 
chlorine.  Iodine  unites  with  bromine  to  form  the  compound  IBr, 
while  a  compound  with  fluorine,  to  which  the  composition  IFB  has 
been  assigned,  is  supposed  to  exist.  None  of  these  compounds  are 
particularly  stable,  and  some  of  them  decompose  easily. 

Exercises.  —  1.  What  impurities  is  commercial  iodine  likely  to 
contain?  In  what  way  does  heating  with  potassium  iodide  (p.  165) 
free  it  from  these? 

2.  Classify  all  the  chemical  actions  in  this  chapter  according  as 
they  belong  to  one  or  other  of  the  ten  kinds  (pp.  124,  163). 

3.  What  are  the  relative  volumes  of  the  gases  in  the  interaction 
of  chlorine  with  hydrogen  bromide  (p.  163),  and  hydrogen  iodide 
(p.  168),  respectively? 


HALOGENS    AS   A    FAMILY  173 

4.  Tabulate,  more  fully  and  specifically  than  is  done  in  the  section 
on  "  The  Halogens  as  a  Family/'  (a)  the  physical  properties,  (b)  the 
chemical  properties,  (c)  the  chemical  relations,  of  the  members  of 
this  group. 

5.  Construct  the  equation  on  p.  164  by  the  use  of  partial  equations 
as  in  the  example  on  p.  160. 

6.  What  are  the  relative  volumes  of  fluorine  and  ozone  in  the 
action  of  the  former  upon  water  (p.  170)? 

7.  What  relative  volumes  of  chlorine  and  iodine  vapor  must  be 
taken  to  make  the  two  chlorides  of  iodine  (p.  172),  respectively? 


-.      : 


CHAPTER  XV 
CHEMICAL  EQUILIBRIUM 

IN  spite  of  its  formidable  title,  this  chapter  will  introduce  nothing 
novel.  Its  purpose  is  to  collect  together  and  organize  more  definitely 
a  number  of  scattered  facts  and  ideas  which  have  already  come  up  in 
various  connections.  On  this  account,  however,  it  will  be  all  the 
more  necessary  for  the  reader  to  refresh  his  remembrance  of  these 
facts  and  ideas  byre-reading  all  pages  to  which  reference  may  be  made. 

Reversible  Actions.  —  In  discussing  the  union  of  hydrogen  and 
iodine  at  445°  (p.  166),  it  was  indicated  that  the  progress  of  the  action 
ceases  while  yet  a  large  amount  of  both  the  substances  necessary  for 
its  maintenance  still  remains  available.  Now  the  materials  left  over 
are  presumably  no  less  capable  of  uniting  than  the  parts  which  have 
already  united.  The  solution  of  this  mystery  lies  in  the  fact  (p.  168) 
that  decomposition  of  the  compound  can  begin  at  180°,  and  therefore 
takes  place  actively  at  445°.  Hence  the  product  of  the  union  must 
begin  to  dissociate,  in  part  at  least,  as  soon  as  any  of  it  is  formed. 
Thus  two  changes,  one  of  which  undoes  the  work  of  the  other,  must 
go  on  simultaneously.  In  consequence  of  this,  neither  can  reach 
completion.  At  445°,  the  system  reaches  equilibrium  when  79  per 
cent  of  hydrogen  iodide  and  21  per  cent  of  the  free  components  are 
present.  As  we  should  expect,  experiment  shows  that  it  makes  no 
difference  whether  we  start  with  the  elements  or  with  the  compound : 
the  proportions  of  the  materials  found  in  the  tube,  after  it  has  been 
heated  for  a  sufficient  length  of  time,  are  in  both  cases  the  same.  A 
general  statement  may  be  founded  on  facts  like  this,  to  the  effect  that 
a  chemical  action  must  remain  more  or  less  incomplete  when  the  reverse 
action  also  takes  place  under  the  same  conditions.  Two  arrows  pointing 
in  opposite  directions  are  used  in  equations  representing  reversible 
changes:*  H2  +  I2  <=±  2HI,  or  2HI  <=>  H2  +  I2. 

*  The  reader  must  avoid  the  idea  that  a  reversible  action  is  one  which  goes 
to  completion,  and  then  runs  back  to  a  certain  extent.  This  conception  would 
be  contrary  to  the  fact,  and  inexplicable  by  the  kinetic  method. 

174 


CHEMICAL   EQUILIBRIUM  175 

The  foregoing  example  of  a  reversible  action,  and  the  following 
examples  which  very  closely  resemble  it,  should  now  be  studied 
attentively  in  connection  with  the  discussion  (for  which  they  furnish 
the  basis)  in  this  and  the  following  sections:  (1)  The  interaction  of 
chlorine  and  water  (p.  115),  which  was  fully  discussed  at  the  time; 
(2)  The  behavior  of  phosphorus  pentachloride  vapor  (p.  146);  (3) 
The  behavior  of  water  vapor  (p.  81),  of  phosphorus  vapor  (p.  147), 
of  sulphur  vapor  (p.  146),  and  of  iodine  vapor  (p.  166). 

When  the  action  is  one  which  is  reversible,  but,  under  the  circum- 
stances being  discussed,  proceeds  far  towards  completion  in  one 
direction,  the  arrow  will  be  modified  to  indicate  this  fact : 

C12  +  H2O  <=5  HC1  +  HC1O  (p.  115). 

When  this  relative  completeness  is  due  to  precipitation  or  volatili- 
zation, the  fact  may  be  indicated  by  vertical  arrows: 

NaCl      +  H2SO4  ^  NaHS04  +  HC1 1  (p.  117). 
NaClJ    +  H2SO4  fc?  NaHS04  +  HC1     (p.  118). 

All  chemical  actions  do  not  belong  to  the  reversible,  incomplete 
class.  Many  proceed  uninterruptedly  to  exhaustion  of  one,  or  all, 
of  the  ingredients.  For  example,  equivalent  amounts  of  magnesium 
and  oxygen  combine  completely,  2Mg  +  O2  — >  2MgO.  Here,  how- 
ever, the  product  is  not  decomposed  even  at  the  white  heat  pro- 
duced by  the  vigor  of  the  union.  Indeed,  magnesium  oxide  cannot 
be  decomposed,  and  the  action  reversed,  at  any  temperature  we  can 
command.  The  other  complete  actions  are  so  because  they  are  likewise 
irreversible. 

Kinetic  Explanation.  —  Restating  these  facts  in  terms  of  the 
kinetic  hypothesis  will  enable  us  to  reason  more  clearly  about  this 
variety  of  chemical  change.  Suppose  we  start  with  the  materials 
represented  on  one  side  of  such  an  equation,  say  the  free  hydrogen 
and  iodine  in  that  on  p.  174.  The  molecules  of  these  materials  will 
encounter  one  another  frequently  in  the  course  of  their  movements. 
In  a  certain  proportion  of  these  collisions  the  chemical  change  will 
take  place.  In  the  earliest  stages  there  will  be  few  of  the  new  kind 
of  molecules  (say  of  hydrogen  iodide),  but,  as  the  action  goes  on, 
these  will  increase  in  quantity.  There  will  be  two  consequences  of 
this.  In  the  first  place  the  parent  materials  will  diminish  in  amount, 
the  collisions  between  their  molecules  will  become  fewer,  and  the 


176  COLLEGE    CHEMISTRY 

speed  of  the  forward  action  will  therefore  become  less  and  less.  In 
the  second  place  the  increase  in  the  number  of  molecules  of  the 
products  will  result  in  more  frequent  collisions  between  them,  in 
more  frequent  occurrence  of  the  chemical  change  which  they  can 
undergo,  and  thus  in  an  increase  in  the  speed  of  the  reverse  action. 
The  forward  action  begins  at  its  maximum  and  decreases  in  speed 
progressively;  the  reverse  action  begins  at  zero  and  increases  in 
speed.  Finally  the  two  speeds  must  become  equal,  and  at  that 
point  perceptible  change  in  the  condition  of  the  whole  must 
cease. 

The  most  immediate  inference  from  this  mode  of  viewing  the 
matter  is,  that  the  apparent  halt  in  the  progress  of  the  action  does 
not  indicate  any  cessation  of  either  chemical  change.  Both  changes 
must  go  on  in  consequence  of  the  continued  encounters  of  the  proper 
molecules.  But  since  the  two  changes  proceed  with  equal  speed  they 
produce  no  alteration  in  the  mass  as  a  whole.  In  fact,  the  final  state 
is  one  of  equilibrium,  and  not  of  rest,  one  of  poise  and  not  of  repose. 
Hence,  chemical  changes  which  are  reversible  lead  to  that  condition 
of  seemingly  suspended  action  which  we  speak  of  as  chemical  equili- 
brium. The  changes  themselves  are  called  reversible,  or,  since  they 
arrive  at  a  state  of  balance  between  opposing  tendencies,  balanced 
actions. 

Chemical  Equilibrium  and  its  Characteristics.  —  The  de- 
tailed discussion  of  the  relations  of  liquid  and  vapor  (pp.  78,  90-92), 
and  of  saturated  solution  and  undissolved  solid  (pp.  101,  105-107), 
has  already  familiarized  us  with  the  term  equilibrium  and  its  signifi- 
cance. By  the  use  of  the  kinetic  hypothesis  we  can,  in  fact,  apply 
sets  of  the  ideas  elaborated  in  these  connections  to  the  discussion 
of  any  kind  of  reversible  phenomena. 

In  particular,  the  reader  will  note  that  the  three  characteristics  of  a 
state  of  equilibrium,  developed  and  illustrated  in  the  case  of  the  physi- 
cal equilibrium  between  a  liquid  and  its  vapor  (p.  91),  apply  also 
to  a  typical  case  of  chemical  equilibrium,  such  as  that  of  hydrogen 
and  iodine  before  us.  Thus: 

1.  There  are  the  two  opposing  tendencies,  which  ultimately  balance 
one  another.  Here  they  are  the  tendency  of  the  hydrogen  and  iodine 
to  produce  hydrogen  iodide,  and  the  tendency  of  the  hydrogen  iodide 
to  reproduce  the  elements  by  its  decomposition.  In  other  words 


CHEMICAL  EQUILIBRIUM  177 

they  are  the  apparent  activity  of  the  hydrogen  and  iodine  interaction, 
and  the  apparent  activity  *  of  the  hydrogen  iodide  decomposition. 

2.  At  equilibrium  the  two  opposing  tendencies  or  activities  are  still 
in  full  operation,  although  their  effects  then  neutralize  one  another. 

3  (and  this  is  the  chief  mark  of  chemical,  as  it  is  of  physical 
equilibrium).  The  system  is  in  a  sensitive  state,  so  that  a  change  in 
the  conditions  (temperature  and  pressure  or  concentration),  even  if 
slight,  produces  a  corresponding  change  in  the  state  of  the  system,  and 
does  this  by  favoring  or  disfavoring  one  of  the  two  opposing  tendencies  or 
apparent  activities.  Such  a  change  is  called  a  displacement  of  the  equili- 
brium, for  the  system  settles  down  in  a  new  state  of  equilibrium  corre- 
sponding to  the  changed  conditions.  Thus,  in  the  present  instance,  a 
change  from  445°,  where  79  per  cent  of  the  hydrogen  and  iodine  are 
in  combination,  to  some  other  temperature,  favors  one  or  other  of 
the  two  opposing  actions  (according  as  the  new  temperature  is  higher 
or  lower  than  before),  and  measurement  at  the  new  temperature 
shows  that  the  proportion  of  the  combined  to  the  free  material  has 
altered.  The  hydrogen  iodide  equilibrium  is  affected  by  changes  in 
concentration  also,  as  we  shall  presently  see  (pp.  179,  183). 

Now,  the  foregoing  facts  show  that  the  key  to  understanding  ap- 
parent chemical  activities,  their  magnitudes,  their  changes,  and  their 
practical  results,  must  lie  in  knowing  how  changes  in  the  conditions 
affect  them.  Hence,  to  the  chemist,  familiarity  with  the  influence  of 
conditions  on  chemical  phenomena  must  be  of  the  greatest  practical 
importance.  We  therefore  address  ourselves  to  the  discussion  of 
this  subject. 

The  "  conditions  "  to  be  considered  are  familiar,  —  temperature, 
and  concentration  or  pressure.  The  "  apparent  activity  of  an 
action  "  which  is  affected  by  these  conditions,  on  the  other  hand,  is 
a  less  easily  specified  thing.  But,  in  cases  of  equilibrium  at  least, 
it  is  accurately  measured  by  the  speed  with  which  the  action  pro- 
ceeds. Thus,  if  the  foregoing  section  be  reexamined,  it  will  be  seen 
that  we  spoke  throughout  of  the  speed,  rather  than  of  the  ten- 
dency or  activity.  So  that  when  we  require  to  consider  some- 

*  We  use  the  term  apparent  activity  for  the  activity  as  we  see  it.  In  the 
same  action  it  varies  with  the  conditions.  The  intrinsic  activity  or  affinity,  on 
the  other  hand,  is  the  absolute  activity  of  the  action  irrespective  of  conditions. 
Its  value  can  be  determined  only  by  eliminating  the  effect  of  conditions,  a 
matter  which  is  too  abstract  for  consideration  here.  The  apparent  activity 
is  the  practical  thing  which  we  observe. 


178  COLLEGE   CHEMISTRY 

thing  more  definite  than  the  activity  we  shall  use  the  speed  of  the 
action. 

Finally,  temperature  and  other  conditions  influence  also  the  activi- 
ties in,  and  therefore  the  speeds  of,  those  actions  which  proceed  to 
completion,  and  are  not  reversible.  Hence,  unless  our  statements 
are  expressly  restricted  to  reversible  actions  and  to  states  of  equili- 
brium, they  apply  to  all  chemical  changes. 

The  Influence  of  Temperature.  —  The  activity  of  chemical 
change,  and  therefore  the  speed  of  all  chemical  changes,  is  increased 
by  raising  the  temperature  and  diminished  by  lowering  it  (cf.  p.  53). 
Thus,  zinc  displaces  hydrogen  more  rapidly  from  hot  than  from  cold 
hydrochloric  acid.  Different  actions  are  affected  in  different  de- 
grees, and  no  simple  rule  accurately  defining  the  effect  can  be  given. 
Roughly  speaking,  however,  a  rise  of  10°  doubles  the  speed  of  every 
action. 

In  a  reversible  change  the  two  opposing  actions  are  different 
actions  and  their  speeds  are  therefore  affected  in  different  degrees  by 
the  same  alteration  in  temperature.  Hence,  when  the  temperature 
is  changed,  the  relative  amount  of  the  two  sets  of  materials  present 
is  altered  and  the  equilibrium  is  displaced.  The  displacements  of 
the  equilibrium  by  raising  or  lowering  the  temperature  were  men- 
tioned in  the  description  of  each  of  the  actions  in  the  reference  list 
in  the  first  section  of  this  chapter.  Thus,  when  phosphorus  penta- 
chloride  is  heated  (p.  146),  the  vapor  is  a  mixture  of  the  pentachloride 
with  the  trichloride  and  free  chlorine:  PC15  +±  PC13  +  C12.  At  200°, 
51.5  per  cent  of  the  material  is  present  as  pentachloride  and  48.5 
per  cent  as  trichloride  and  chlorine.  Raising  the  temperature  to 
250°  changes  the  proportions  to  20  per  cent  and  80  per  cent,  respec- 
tively. At  300°  only  3  per  cent  of  the  pentachloride  remains. 
Evidently,  here,  raising  the  temperature  favors  the  decomposition  of 
the  pentachloride,  and  therefore  increases  the  speed  of  its  dissocia- 
tion more  than  it  does  the  speed  of  the  re-union  of  the  trichloride 
and  chlorine. 

The  Influence  of  Concentration.  —  Leaving,  now,  the  tempera- 
ture out  of  consideration,  or  considering  it  to  be  constant,  we  take 
up  the  influence  of  concentration  upon  apparent  activity.  We  have 
seen  (p.  175)  that  the  speed  of  a  chemical  change  is  determined  by 


CHEMICAL  EQUILIBRIUM  179 

the  frequency  with  which  the  molecules  of  the  necessary  substances 
encounter  one  another.  The  frequency  of  the  encounters  amongst 
a  given  set  of  molecules,  resulting  in  a  definite  chemical  change,  will 
in  turn  evidently  depend  entirely  upon  the  degree  to  which  the  mole- 
cules are  concentrated  in  each  other's  neighborhood.  Larger 
amounts  of  one  of  the  materials,  for  example,  will  not  result  in  more 
rapid  chemical  action,  if  the  larger  amount  of  material  is  also  scat- 
tered through  a  larger  space.  Chemical  changes,  therefore,  are  not 
accelerated  by  increasing  the  mere  quantity  of  any  ingredient,  but 
only  by  increasing  the  concentration  of  its  molecules.  Thus,  a  large 
amount  of  hydrochloric  acid  with  a  piece  of  zinc  will  generate  hydro- 
gen no  faster  than  a  smaller  amount.  But  substitution  of  more  con- 
centrated acid  will  instantly  increase  the  speed  of  the  action.  In 
the  second  case,  the  number  of  molecules  of  the  acid  reaching  the 
zinc  per  second  is  greater,  and  this  action,  being  non-reversible, 
proceeds  more  rapidly  to  complete  consumption  of  the  zinc.  With 
a  reversible  action  the  effect  on  the  speed  is  the  same,  excepting  that 
the  continued  activity  of  the  reverse  action  prevents  the  direct  one 
from  reaching  completion.  Thus,  if,  in  the  action  of  hydrogen  upon 
iodine,  we  introduce  into  the  same  space  an  extra  amount  of  hydrogen, 
this  facilitates  the  formation  of  hydrogen  iodide  by  increasing  the 
possibilities  of  encounter  between  hydrogen  and  iodine.  At  the 
same  time  it  does  not  affect  (cf.  p.  60)  the  number  of  encounters  in 
a  given  time  of  hydrogen  iodide  molecules  with  one  another  which 
result  in  the  reverse  transformation.  The  proportion  of  hydrogen 
iodide  formed,  therefore,  from  a  given  amount  of  iodine  will  be 
greater,  although  the  total  possible  (by  complete  consumption  of 
the  materials)  has  not  been  altered,  since  the  quantity  of  one  ingre- 
dient only  has  been  increased.  The  introduction  of  an  excess  of 
iodine  would  have  had  precisely  the  same  effect. 

It  is  easy  to  illustrate  this  experimentally.  If  ferric  chloride 
and  ammonium  thiocyanate  are  mixed  in  aqueous  solution,  a  liquid 
containing  the  soluble,  blood-red  ferric  thiocyanate  is  produced.  The 
compound  radicals  are  (NH4)  and  (CNS),  and  the  action  is  a  simple 
double  decomposition: 

FeCl3  +  3NH4CNS  «=»  Fe(CNS)3  +  3NH4C1. 

The  action  is  a  reversible  one,  and  the  mixture  is  homogeneous,  i.e. 
there  is  no  precipitation.  Now,  if  the  two  just-named  salts  are 


180  COLLEGE   CHEMISTRY 

mixed  in  very  dilute  solution  in  the  proportions  required  by  the 
equation,  say  by  adding  20  c.c.  of  a  decinormal  solution  (p.  99)  of 
each  salt  to  several  liters  of  water,  a  pale-reddish  solution  is  obtained. 
When  this  is  divided  into  four  parts,  and  one  is  kept  for  reference, 
the  addition  of  a  little  of  a  concentrated  solution  of  ferric  chloride 
to  one  jar,  and  of  ammonium  thiocyanate  to  another,  will  be  found 
to  deepen  the  color  by  producing  more  of  the  ferric  thiocyanate.  On 
the  other  hand,  mixing  a  few  drops  of  concentrated  ammonium 
chloride  solution  with  the  fourth  portion  will  be  found  to  remove 
the  color  almost  entirely,  on  account  of  its  influence  in  accelerating 
the  backward  change. 

The  general  principle  discussed  and  illustrated  in  this  section  may 
be  called  the  law  of  molecular  concentration,  and  may  be  stated  as 
follows:  In  every  chemical  change  the  apparent  activity,  and  therefore 
the  speed  of  the  action,  at  any  moment,  is  proportional  to  the  molecular 
concentration  for  the  time  being  of  each  interacting  substance.  This 
holds  whether  the  action  is  reversible  or  not. 

We  shall  next  give  a  more  precise,  semi-mathematical  formulation 
of  this  law,  as  this  formulation  will  be  of  use  later,*  and  then  proceed 
to  illustrate  the  application  of  the  law,  by  showing  how  it  explains 
large  classes  of  actions  of  which  we  have  already  encountered  many 
examples. 

*  Formulation  of  the  Law  of  Molecular  Concentration.  —  The 

mathematical  formulation  of  the  law  describing  the  influence  of  the 
concentration  of  the  molecules  of  each  participating  substance  upon 
the  speed  of  the  action,  and  therefore  upon  the  apparent  activity 
of  the  action,  is  extremely  simple.  When  the  actual  concentrations 
of  the  molecules  are  specified  (in  moles  (pp.  99,  129)  per  liter),  and 
the  speed  is  suitably  expressed  (in  moles  transformed  per  minute  or 
per  hour),  we  find  that  the  speed  is  proportional  to  the  concentration 
of  each  molecule  appearing  in  the  molecular  equation  for  the  action. 
Thus  in  the  interaction  of  hydrogen  and  iodine,  if  [H2]  and  [I2]  repre- 
sent the  concentrations  of  the  molecules  H2  and  I2,  and  &  is  a  con- 
stant, and  S  is  the  speed,  then: 

[H2]  X  [I2]  X  k  =  S. 

*  This  mathematical  formulation  of  the  law  is  not  required,  or  referred 
to  in  the  sections  which  follow.     The  section  may  therefore  be  omitted  for 
the  present  and  taken  up  in  connection  with  Chap,  xxxiv. 


CHEMICAL  EQUILIBRIUM  181 

Again,  for  the  dissociation  of  phosphorus  pentachloride  vapor  into 
phosphorus  trichloride  and  chlorine  (p.  146):  PC16— >  PC13  +  C12,  if 
[PC15]  represent  the  concentration  of  the  PC15  molecules,  kl  is  a  con- 
stant, and  Si  is  the  speed  of  decomposition: 

[raj  xv-  s* 

Similarly,  for  the  reverse  action:  PC13  +  C12  ->  PC16,  if  [PC13]  and 
[C12]  stand  for  the  molecular  concentrations  of  these  substances: 

[PC13]  X  [C1J  X  k2  =  Sy 

The  constant  has  a  different  value  in  each  separate  action.  It 
includes  the  value  of  the  intrinsic  affinity  or  activity  of  the  sub- 
stances, and  the  catalytic  effect  (p.  54),  if  any,  of  the  materials 
present. 

The  foregoing  plan  may  be  used  further  to  formulate  the  condition 
for  chemical  equilibrium.  As  we  have  seen  (p.  176),  a  characteristic 
of  a  system  in  chemical  equilibrium  is  that  the  apparent  activities  of 
the  opposing  actions  balance  one  another,  and  therefore  the  speeds 
of  the  forward  and  reverse  actions  have  become  equal.  If,  then, 
[PCl6]eqm.»  [PCl3]eqm.,  and  [Cl2]eqm.  now  represent  the  molecular  con- 
centrations when  the  system  has  reached  equilibrium,  then,  since 
the  speeds  are  equal : 

[PCl3]eqm.   X   [Cljeqm.   X   &2   =   [PClJeqm.    X  k, 

[pcyeqm.  x  [ci2]eqm.  j±  _  constant 

[PCl5]eqm.  &2 

In  words,  this  means  that  if  we  change  the  amount  of  the  penta- 
chloride placed  in  the  vessel,  or  if  we  use  amounts  of  chlorine  and 
trichloride  which  are  not  equivalent,  the  numerical  value  at  equili- 
brium of  each  concentration  ([PC13]  etc.)  will,  of  course,  be  different, 
but  the  product  of  the  concentrations  of  trichloride  and  chlorine, 
divided  by  the  concentration  of  the  pentachloride,  will  always  give 
the  same  numerical  value  for  the  constant  at  the  same  temperature. 
This  numerical  value  is  called  the  equilibrium  constant. 

Applications:  The  Forward  Action.  Homogeneous  and 
Inhomoyeneous  Systems.  —  While  there  are  all  degrees  of  speed 
in  chemical  actions,  yet  in  practice  we  quickly  distinguish  two  differ- 
ent classes.  There  is  a  class  of  actions  of  which  most  examples  are 
almost  instantaneously  accomplished,  and  a  class  in  which,  fre- 


182  COLLEGE   CHEMISTRY 

quently,  the  operation  takes  minutes  or  even  hours.  The  classes 
overlap,  but,  in  a  general  way,  the  following  distinction  may  be 
made. 

To  the  former,  speedy  class  belong  the  explosion  of  hydrogen  and 
oxygen  or  other  gaseous  mixtures,  and  the  interactions  when  solu- 
tions are  mixed,  as  in  precipitations.  In  view  of  the  foregoing 
explanations,  we  perceive  that  the  rapid  accomplishment  of  such 
actions  is  due,  not  so  much  to  any  especially  great  intrinsic  affinity,  as 
to  the  homogeneous  state  of  mixture  of  the  interacting  materials.  This, 
of  course,  is  a  purely  physical,  and  not  a  chemical  motive  for  speedy 
interaction.'  In  intimate  mixtures,  every  molecule  has  an  equal 
opportunity  freely  to  encounter  every  other  molecule  and  there  is 
therefore  no  mechanical  impediment  to  the  operation  of  the  affinities 
of  the  substances.  Hence  the  apparent  activity  is  great. 

To  the  second  class,  comprising  the  slower  actions,  belong  cases 
like  the  interaction  of  a  piece  of  zinc  with  hydrochloric  acid,  or  of 
manganese  dioxide  (p.  Ill)  with  the  same  acid,  whereby  hydrogen 
and  chlorine,  respectively,  are  slowly  evolved,  and  the  solid  is  grad- 
ually consumed.  Here  the  hindrance  is  evidently  the  fact  that  the 
interacting  substances  are  not  intimately  mixed.  In  the  slow 
actions,  the  system  is  inhomogeneous.  Pulverizing  the  solid  before  use 
will  increase  the  speed,  indeed,  but  will  not  transfer  the  action  to 
the  rapid  class.  It  is  chiefly  the  dissolved  part  of  the  substance  which 
interacts,  for  chemical  action  takes  place  between  molecules,  and 
only  the  dissolved  part  is  disintegrated  in  such  a  way  that  the 
molecules  are  readily  accessible.  Thus,  the  action  is  held  back  by 
continual  waiting  for  the  slow  replenishment,  from  the  "  insoluble  " 
solid,  of  the  supply  of  dissolved  molecules.  In  the  cases  cited,  the 
restraining  influence  of  the  dissolving  process,  which  is  part  of  the 
whole  phenomenon,  may  be  formulated  thus : 

Zn(solid)  ±?  Zn(dslvd)  +  2HC1  -»  ZnCl2  +  H2. 
MnO,  (solid)  +5  MnO2(dslvd)  +  4HC1  ->  MnCl2  +  2H2O  +  C12. 

Here,  again,  the  mechanical  details,  depending  on  physical  proper- 
ties, have  more  to  do  with  the  progress  of  the  action  than  has  the 
chemical  affinity.  In  terms  of  the  law  of  concentration,  the  action 
is  slow,  and  the  apparent  activity  small,  because  the  concentration 
of  the  acting  molecules  of  one  of  the  substances  is  very  small,  and 
cannot  be  increased  because  of  low  solubility  (cf.  p.  162,  first  line). 


CHEMICAL  EQUILIBRIUM  183 

Applications  :  TJie  Reverse  Action.  Displacement  of  Equili- 
bria. —  We  have  just  seen  (p.  179)  that  one  way  in  which  a  reversible 
action  may  be  forced  nearer  to  completion,  in  one  direction  or  the 
other,  is  the  introduction  of  an  excess  of  one  of  the  ingredients  con- 
tributing to  the  forward  action.  This  method  of  displacing  the 
equilibrium  point,  however,  cannot  be  very  effective,  unless  it  is 
possible  to  introduce  an  exceedingly  large  excess  of  the  selected 
ingredient  in  a  high  degree  of  molecular  concentration,  since  this 
operation  does  not  in  any  way  effect  or,  in  particular,  restrain  the 
reverse  action  which  is  continually  undoing  the  work  of  the  forward 
one.  A  much  more  effective  means  of  furthering  the  desired  direction 
of  such  actions  is  found,  therefore,  in  the  restraint  or  practical  annulment 
of  the  reverse  action.  A  good  way  of  accomplishing  this  is  to  allow 
the  products  of  the  direct  action  to  separate  into  an  inhomogeneous 
mixture.  Any  agency  which  could  remove  the  free  iodine  vapor  as 
fast  as  it  was  formed  in  the  decomposition  of  hydrogen  iodide,  for 
example,  would  entirely  stop  the  reproduction  of  the  compound,  and 
so  would  enable  the  dissociation  (2HI  — >  H2  +  I2)  to  run  to  com- 
pletion. 

This  might  be  realized  by  causing  one  end  of  a  sealed  tube  charged 
with  hydrogen  and  iodine,  after  the  contents  had  settled  down  to  a 
condition  of  equilibrium,  to  project 
from  the  bath  in  which  the  whole 
had  been  kept  at  445°    (Fig.   49, 
which  is  simply  diagrammatic) .    By 
cooling  this  end,  a  large  part  of  the 
21  per  cent  of  free  iodine  would 


quickly  be  condensed  in  it  to  the  FIG.  49. 

solid    form,    while    the    hydrogen 

would  remain  gaseous.  In  other  words  the  concentration  of  the 
free  iodine  would  be  greatly  reduced.  In  fact,  only  the  trace  of 
vapor  which  cold  iodine  gives  would  then  be  available  to  interact 
with  the  hydrogen,  and  reproduce  hydrogen  iodide.  Meanwhile  the 
decomposition  of  the  latter  would  go  on,  and  thus,  eventually,  almost 
all  the  iodine  would  be  found  free  in  one  end  of  the  tube,  and  the 
hydrogen,  all  free  likewise,  would  occupy  the  rest.  By  this  purely 
mechanical  adjustment  the  chemical  change  would  therefore  be 
carried  from  21  per  cent  completion  to  almost  absolute  completion: 

2HI  *=;  H2  +  I2(vapor)  fc?  I, (solid). 


184  COLLEGE   CHEMISTRY 

If,  on  the  other  hand,  arrangements  were  made  to  have  powdered 
marble,  in  a  sealed  bulb  of  thin  glass,  inclosed  in  the  tube,  we  might 
imagine  the  very  opposite  effect  of  the  above  to  be  produced.  The 
breaking  of  the  bulb  of  marble,  when  equilibrium  had  been  reached, 
would  provide  means  for  the  removal  of  all  the  hydrogen  iodide,* 
while  the  hydrogen  and  iodine  would  still  be  gaseous.  Thus,  the 
compound  (HI)  having  been  reduced  in  concentration  to  the  point 
of  being  removed  entirely,  there  would  be  no  reverse  action  to  com- 
pensate for  the  union  of  the  elements.  The  whole  material  would, 
therefore,  soon  have  passed  through  the  form  HI.  Hence,  by 
another  mechanical  arrangement,  an  action  which  ordinarily  could 
progress  to  only  79  per  cent  would  be  turned  into  a  complete  one: 
H2  +  I2  fc?  2HI  (+  CaCO3  -»  CaI2  +  H20  +  C02). 

In  every-day  chemical  work,  since  our  object  is  usually  to  prepare 
some  one  substance,  chemists  either  avoid  chemical  changes  which 
are  notably  reversible,  or  adjust  the  conditions,  as  is  done  in  the 
foregoing  illustrations,  so  that  the  reverse  of  the  action  which  they 
desire  is  prevented.  In  consequence  of  this,  when  carrying  out  the 
directions  for  making  familiar  preparations,  the  fact  that  such  actions 
are  reversible  at  all  very  readily  escapes  our  notice.  Arranging 
the  conditions  so  that  the  separation  of  a  solid  body  by  pre- 
cipitation, or  the  liberation  of  a  gas,  takes  place,  are  the  two  com- 
monest ways  of  rendering  a  reversible  action  complete.  Excellent 
examples  of  both  of  these  are  furnished  by  the  chemical  change 
used  in  producing  hydrogen  chloride  by  the  inter-action  of  salt  and 
sulphuric  acid,  the  full  discussion  of  which  (p.  118)  should  now  be 
studied  attentively  in  the  light  of  these  explanations  (p.  24,  footnote). 

The  reader  will  find  in  Erin's  process  (p.  46),  where  the  concen- 
tration of  the  oxygen  is  changed  by  the  use  of  pumps,  another 
exemplification  of  the  principles  explained  in  this  chapter.  The 
behavior  of  hydrates  (p.  83)  furnishes  a  whole  class  of  illustrations. 

Exercises.  —  1.  Explain  the  completeness  of  the  action  by 
which  hydrogen  chloride  and  water,  respectively,  are  formed  by 
direct  union  of  the  elements. 

*  The  hydrogen  iodide  would  be  destroyed  by  interaction  with  the  marble: 

2HI  +  CaCO3  — >  CaI2  +  CO2  +  H2O. 

The  calcium  iodide  is  a  solid.  The  two  gases,  carbon  dioxide  and  water  vapor, 
are  here  assumed  not  to  interact  with  hydrogen  or  with  iodine,  and  would 
not,  therefore,  interfere  with  the  formation  of  fresh  hydrogen  iodide. 


CHEMICAL   EQUILIBRIUM  185 

2.  Explain  the  completeness  of  the  action  by  which  silver  chloride 
(p.  8)  is  formed. 

3.  Explain   why   the   decomposition   of   potassium   chlorate   is 
complete. 

4.  In  view  of  the  statement  on  p.  45,  explain  why  mercuric  oxide 
is  completely  decomposed  by  heating.     Point  out  the  resemblance 
between  this  experiment  and  the  imaginary  one  illustrated  in  Fig. 
49  (p.  183). 

5.  Why  can  magnetic  oxide  of  iron  be  reduced  completely  by  a 
stream  of  hydrogen  (p.  75),  and   iron  oxidized   completely  by  a 
current  of  steam  (p.  67)  ? 

6.  What  actions  in  Chap,  xiv  are  complete  for  the  same  reason 
that  the  action  of  sulphuric  acid  on  salt  (pp.  161,  167,  170)  is  so? 

7.  With  the  phosphorus  pentachloride  system,  say  at  250°,  what 
effect  would  suddenly  enlarging  the  space  containing  a  given  amount 
of  the  vapor  produce?     What  would  be  the  effect  of  diminishing  the 
space?    What  would  be  the  effect  of  introducing  additional  chlorine 
into  the  same  space  (p.  179)? 

8.  By  what  practical  means  could  the  degree  of  dissociation  of 
sulphur  vapor  (S8)  be  reduced,  without  changing  the  temperature 
(p.  146)? 


CHAPTER  XVI 
OXIDES    AND    OXYGEN    ACIDS    OF    THE    HALOGENS 

THE  chief  subjects  of  practical  importance  touched  upon  in  this 
chapter  are  connected  with  bleaching  powder  (CaCl(OCl)),  and  potas- 
sium chlorate  (KC1O3)  and  perchlorate  (KC1O4).  Hence  our  atten- 
tion will  be  largely  directed  to  the  modes  of  making  these  substances 
and  to  their  relations  to  one  another.  Incidentally,  we  shall  encoun- 
ter many  actions  of  a  complex  and,  to  us,  more  or  less  novel  kind. 

Compounds  of  Chlorine  Containing  Oxygen.  —  The  following 
are  the  names  and  formulae  of  the  substances: 

HC10  Hypochlorous  acid,  C12O  Hypochlorous  anhydride, 

[HC1O2]  Chlorous  acid,  

C102  Chlorine  dioxide, 

HC1O3  Chloric  acid,  

HC1O4  Perchloric  acid,  C1207  Perchloric  anhydride. 

There  are  also  compounds  of  metals  with  the  negative  radicals  of 
these  acids.  Of  this  nature  are  the  three  substances  mentioned  in 
the  first  paragraph.  Chlorous  acid  is  itself  unknowrn,  but  potassium 
chlorite  (KC1O2)  and  some  other  derivatives  have  been  made. 

The  two  anhydrides  (p.  51),  when  brought  into  contact  with 
water,  combine  with  it  to  form  the  acids  opposite  which  they  stand 
in  the  table.  Chlorine  dioxide  (q.v.),  however,  is  not  related  to  any 
one  acid  in  this  way. 

All  these  compounds  differ  from  most  that  we  have  hitherto  dis- 
cussed, inasmuch  as  not  one  of  them  can  be  made  by  direct  union  of 
the  simple  substances. 

Nomenclature  of  Acids  and  Salts»  —  When  several  compounds 
closely  related  in  composition,  like  the  above  acids,  are  known,  a 
systematic  method  of  naming  them  is  used.  The  terminations  -ous 
and  -ic  indicate  smaller  and  larger  proportions  of  oxygen  respectively 

186 


OXIDES   AND   OXYGEN   ACIDS  OF    CHLORINE  187 

(cf.  p.  51).  For  compounds  below  or  above  those  two  in  their 
degree  of  oxidation,  the  prefixes  hypo-  and  per-  are  employed. 

When  the  negative  radicals  (p.  64)  of  the  acids  are  combined 
with  metals,  the  compounds  are  spoken  of  as  salts  of  the  respective 
acids.  Thus,  KC103  is  described  as  the  potassium  salt  of  chloric 
acid.  The  specific  names  for  these  salts  are  distinguished  by  ter- 
minations corresponding  to  those  of  the  acids: 

KC1O   Potassium  hypochlorite,  HC1O   Hypochlorous  acid, 

KC102  Potassium  chlorite,  HC1O2  Chlorous  acid, 

KC1O3  Potassium  chlorate,  HC1O3  Chloric  acid, 

KC1O4  Potassium  perchlorate.  HC1O4  Perchloric  acid. 

The  termination  -ite  corresponds  to  -ous,  -ate  to  -ic.  This  principle 
is  applied  systematically,  so  that  the  salts  of  sulphuric  and  sulphur- 
ous acids,  for  example,  are  called  sulphates  and  sulphites  respectively 
(cf.  p.  71). 

Compounds  containing  two  elements  only  receive  the  termina- 
tion -ide.  Thus,  KCL  is  potassium  chloride,  FeS  is  ferrous  sulphide. 

Salts  and  Double  Decomposition.  —  We  have  just  been  using 
the  word  salt  in  a  general  sense  (cf.  p.  157).  It  is  the  class  name  for 
a  set  of  substances  which  includes  common  salt  or  sodium  chloride 
(NaCl),  potassium  nitrate  (KNO3),  sodium  sulphate  (Na2S04),  silver 
chloride  (AgCl),  potassium  chlorate  (KC103),  etc.  The  majority  of 
the  substances  used  in  elementary  chemistry  belong  to  this  class. 
They  receive  the  name  because  in  certain  important  chemical  respects 
they  behave  like  common  salt.  For  example,  when  sodium  chloride 
was  treated  with  sulphuric  acid  (p.  117)  or  phosphoric  acid  (p.  118), 
an  exchange  of  radicals  took  place.  An  action  of  the  same  type  was 
that  of  sodium  chloride  and  silver  nitrate  in  aqueous  solution  (p.  8). 
Here  we  have  two  salts  interacting,  instead  of  an  acid  and  a  salt, 
and  the  interchange  of  radicals  is  exactly  similar: 

NaCl  +  H2(SO4)  <=>  NaH(SO4)  +  HC1 1 
NaCl  +  AgNO3    <=>  NaNO3  +  AgCl  J. 

Now  salts  in  general  behave  in  these  respects  in  the  same  way  as 
does  common  salt.  They  interact  with  acids  or  other  salts,  particularly 
in  solution,  in  such  a  way  that  an  exchange  of  radicals  takes  place.  In 
the  first  case,  a  salt  and  an  acid,  and  in  the  second  case  two  salts, 


188  COLLEGE   CHEMISTRY 

are  produced.  These  actions  are  all  reversible.  Acids  differ  thus 
from  salts  only  in  the  fact  that  one  of  their  radicals  is  hydrogen. 
Hence  they  are  frequently  called  hydrogen  chloride,  hydrogen  sul- 
phate (H2SO4),  and  so  forth.  It  may  be  added  that  bases  (p.  81), 
like  potassium  hydroxide  (KOH),  interact  reversibly  with  salts  and 
acids,  exchanging  radicals  after  the  same  fashion  (e.g.  pp.  123, 
189,  191). 

All  salts  are  named  according  to  the  radicals  which  they  contain. 
Thus,  all  containing  S04  are  sulphates.  Conversely,  when  the  name 
of  a  salt  is  given,  the  formula  can  be  written  down  at  once.  In  doing 
this,  however,  regard  must  be  had  to  the  valence  of  the  radicals 
(p.  70). 

In  view  of  the  reversibility  of  most  of  the  interactions  of  salts, 
acids,  and  bases,  we  encounter  completed  changes  chiefly  when  pre- 
cipitation occurs,  or  when  one  product  is  volatile  (p.  184) .  If  neither 
of  the  products  formed  by  the  exchange  of  radicals  is  insoluble,  the 
reversibility  of  the  action  prevents  our  obtaining  anything  but  a 
mixture.  Only  those  double  decompositions  which  involve  more  or 
less  insoluble  or  volatile  substances  are  thus  of  use  for  preparing 
salts.  The  action  of  sodium  chloride  on  silver  nitrate  is  an  example. 
The  silver  chloride  is  almost  completely  insoluble,  while  the  sodium 
nitrate  produced  by  the  change  remains  dissolved.  By  filtration  we 
obtain  the  silver  chloride  as  a  powder,  while  the  evaporation  of  the 
filtrate  gives  us  the  soluble  product.  This  sort  of  action  can  be  used, 
therefore,  either  for  the  preparation  of  a  soluble  or  an  insoluble  substance. 
If  the  problem  is  to  make  a  soluble  product,  then  we  must  arrange  an 
action  between  two  substances,  each  containing  one  of  the  two 
required  radicals,  and  possessing  two  other  radicals,  which,  when 
united,  give  an  insoluble  body.  This  plan  is  illustrated  frequently 
in  what  follows  (e.g.  pp.  195,  197). 

Preparation  and  Properties  of  Hypochlorites.  —  Since  none 
of  the  acids  in  our  list  can  be  made  directly  from  their  elements, we 
generally  have  to  prepare,  first,  the  corresponding  salt.  From  the 
salt,  by  double  decomposition,  the  acid  is  then  secured.  Hence,  in 
each  case,  the  salts  will  be  discussed  first. 

Chlorine  interacts  slightly  with  water  (p.  115),  producing  small 
quantities  of  hydrogen  chloride  and  hypochlorous  acid  (equation  (1), 
below).  The  action  is  very  strongly  reversible.  That  is  to  say,  since 


OXIDES  AND   OXYGEN   ACIDS   OF   CHLORINE  189 

the  last  two  substances  interact  very  vigorously  to  reproduce  chlorine 
and  water,  the  direct  action  does  not  make  much  progress. 

When,  however,  some  substance  which  can  interact  with  one  or 
both  of  these  products  is  added  to  the  solution  of  chlorine,  or  when 
chlorine  gas  is  simply  passed  into  an  aqueous  solution  of  such  a 
substance,  displacement  of  the  equilibrium  point  at  once  occurs 
(p.  183).  Now  potassium  hydroxide  is  a  suitable  substance.  It 
interacts  almost  completely  in  solution  with  both  the  products  of 
this  action,  producing  potassium  chloride  (2)  and  potassium  hypo- 
chlorite  KOC1  (3),  according  to  the  last  two  of  the  following  equations: 

C12       +  H2O    ±=5  HC1     +  HOC1,  (1) 

HC1     +  KOH  -» KC1     +  H20,  (2) 

HOC1  +  KOH  -»  KOC1  +  H2O.*  (3) 

Thus,  omitting  the  water,  which  appears  both  among  products  and 
initial  substances  and  in  any  case  is  present  in  large  excess  as  a 
solvent,  and  omitting  also  the  two  acids,  which  are  used  up  as 
quickly  as  they  are  produced  by  equation  (1)  and  are  not  amongst 
the  actual  products,  we  get,  by  addition  of  the  three  partial  equations 
(cf.  p.  159),  the  final  equation: 

C12  +  2KOH  ->  KC1  +  KOC1  +  H2O. 

This  sort  of  action  does  not  give  pure  potassium  hypochlorite,  but 
for  some  purposes  the  presence  of  potassium  chloride  in  the  solution 
is  not  objectionable. 

Bleaching  powder,  CaCl(OCl),  is  manufactured  on  a  large  scale  by 
an  action  exactly  like  the  above.  The  neutralization*  of  a  molecule 
of  each  of  the  two  acids,  however,  can  be  accomplished  by  a  single 
molecule  of  slaked  lime  (calcium  hydroxide,  Ca(OH)2),  since  the 
latter  contains  two  hydroxyl  (OH)  groups.  The  lime  can  be  applied, 
either  in  the  dry  form  or  mixed  with  some  water  as  a  paste.  The 
separate  actions  and  final  equation  are  as  follows: 

C12  +   H20     ±^HC1          +HOC1  (1) 

r«o  /  On.  +   HC1      — »pt   /  Cl       ,   QTT  f\  /9\ 

Ca\OH+  HOC1  -*Ca\OCl  ' 

C12     +Ca(OH)2->CaCl(OCl)  +  H2O 

*  The  double  decomposition  of  potassium  hydroxide,  or  any  other  base, 
with  any  acid  to  produce  a  salt  and  water,  is  called  neutralization  (cf. 
p.  188). 


190  COLLEGE   CHEMISTRY 

Bleaching  powder  (q.v.)  is  a  salt  of  calcium  involving  two  different 
acid  radicals  (a  mixed  salt).  This  condition,  again,  does  not  inter- 
fere with  the  application  of  the  substance  commercially.  A  method 
of  obtaining  pure  hypochlorites,  however,  will  be  found  below. 

Three  chemical  properties  of  hypochlorites  deserve  mention:  They 
change  into  chlorates  (q.v.)  when  heated.  They  may  also  give  off 
oxygen,  2CaCl(OCl)  ->  2CaCl2  +  02.  Although  this  decomposition 
is  slow  in  cold  solutions  of  hypochlorites,  or  when  they  are  preserved 
in  the  dry  form,  it  may  be  hastened  by  means  of  catalytic  agents. 
The  addition  of  a  little  cobalt  hydroxide  (q.v.)  to  bleaching  powder 
solution  causes  rapid  evolution  of  oxygen.  Finally,  hypochlorites 
interact  with  acids  by  double  decomposition  (cf.  p.  187)  to  give 
hypochlorous  acid  (see  below).  It  is  for  the  purpose  of  getting  this 
acid,  which  is  a  powerful  bleaching  agent,  that  the  hypochlorites 
are  manufactured.  The  salts,  such  as  bleaching  powder,  can  be 
stored  and  transported  easily,  while  the  acid  itself  will  not  keep, 
except  when  largely  diluted,  and  it  consequently  cannot  be  handled 
conveniently. 

Preparation  of  Hypochlorous  Acid  :  Hypochlorous  Anhy- 
dride. —  1.  The  common  method  of  obtaining  the  acid,  HOC1,  is  by 
double  decomposition,  using  a  hypochlorite  with  some  other  acid 
(p.  187).  When  only  a  mixture  of  a  chloride  and  a  hypochlorite, 
such  as  is  produced  by  the  action  of  chlorine  on  a  base  (p.  189),  is 
available,  we  have,  simultaneously,  the  two  reversible  actions: 

(KOC1  +  HN03  fc?  KN03  +  HOC1, 
JKC1     +  HNO3  <=±  KNO3  +  HC1. 

But  hypochlorous  acid  is  a  feeble  acid,  while  hydrochloric  acid  is  an 
active  one,  so  that  in  the  upper  action  the  reversing  tendency  is  very 
slight,  while  in  the  lower  it  is  vigorous.  Hence,  by  adding  nitric 
acid,  in  amount  barely  sufficient  for  the  liberation  of  the  hypochlor- 
ous acid  alone,  and  doing  this  in  a  very  dilute  solution,  the  object  is 
attained.  The  potassium  chloride  is  hardly  affected.  By  gently 
warming  the  liquid  a  dilute  solution  of  hypochlorous  acid  can  be 
distilled  off. 

2.  When  chlorine  is  passed  into  water  holding  chalk  in  suspension, 
only  the  hydrochloric  acid,  produced  by  the  interaction  of  the  chlo- 
rine and  water  (equation  1,  p.  189),  acts  upon  the  chalk.  The 


OXIDES  AND  OXYGEN  ACIDS   OF   CHLORINE  191 

hypochlorous  acid  is  too  feeble  to  interact  and  can  be  obtained  by 
subsequent  distillation: 

2C12  +  CaCO3  +  H2O  -»  CaCl2  +  CO2t  +  2HOC1. 

3.   A  third  method  is  by  addition  of  water  to  the  anhydride. 

The  anhydride  of  hypochlorous  acid  (chlorine  monoxide  C12O)  may 
be  obtained  by  passing  chlorine  gas  over  precipitated  mercuric  oxide. 
Each  of  the  constituents  of  the  oxide  combines  with  chlorine: 

HgO  +  2C12  ->  HgCl2  +  C120. 

The  mercuric  chloride  then  unites  with  another  formula-weight  of  the 
mercuric  oxide  to  form  a  basic  mercuric  chloride  HgO,  HgCl2,  which 
remains  in  the  tube.  The  chlorine  monoxide  is  a  brownish-yellow 
gas.  When  slightly  warmed  it  decomposes  into  its  constituents 
with  explosion.  The  gas  dissolves  in  water  very  easily  (200  :  1, 
by  vol.).  The  yellow  solution  of  hypochlorous  acid  which  results: 

C12O  +  H2O  t=?  2HOC1, 

has  a  strong  odor  of  chlorine  monoxide.  The  combination  is  rever- 
sible, and,  especially  when  the  liquid  is  warm,  a  little  of  the  gas 
escapes. 

Properties  of  Hypochlorous  Acid.  —  1.  Hypochlorous  acid 
cannot  be  made,  excepting  in  solution,  or  kept,  excepting  in  dilute 
solution.  This  is  in  consequence  of  its  tendency  to  decompose  in 
three  different  ways,  one  of  which,  the  liberation  of  the  anhydride, 
has  just  been  mentioned  (see  3  and  4  below). 

2.  As  an  acid  it  neutralizes  (p.  189)  active  bases,  giving  hypo- 
chlorites  in  a  pure  condition:  NaOH  +  HOC1  -»  NaOCl  +  H2O. 

3.  If  the  solution  is  concentrated,  much  of  the  hypochlorous  acid 
changes  gradually  into  chloric  acid  and  hydrogen  chloride.     This 
occurs  even  in  the  dark:  3HOC1  -»  HC1O3  +  2HC1. 

4.  When  the  solution  is  warmed,  but  more  especially  when  it  is 
exposed  to  sunlight,  oxygen  is  evolved  rapidly. 

2HOC1  -*  2HC1  +  O2. 

This  decomposition  always  takes  place  in  sunlight,  whether  the  acid 
is  present  alone  in  the  water,  or  along  with  other  substances.  Hence, 
the  solution  of  chlorine  in  water  (pp.  115,  189),  which  contains  a 
small  amount  of  hypochlorous  acid,  on  being  exposed  to  bright 


192  COLLEGE   CHEMISTRY 

sunlight,  gives  off  bubbles  of  oxygen  rapidly.  This  decomposition, 
since  it  removes  one  of  the  interacting  substances  in  the  reverse 
action,  C12  +  H2O  ±5  HC1  +  HOC1,  enables  the  interaction  of  chlo- 
rine and  water  to  go  on  to  completion.  Consequently,  the  final 
liquid  contains  hydrochloric  acid  and  water.  Leaving  out  the 
intermediate  steps  again,  the  action  appears,  therefore,  to  be  simply 
a  decomposition  of  water  by  chlorine:  2C12  +  2H2O  — »  4HC1  +  O2. 

5.  In  consequence  of  the  ease  with  which  it  gives  up  oxygen, 
hypochlorous  acid  is  a  strong  oxidizing  agent  (see  below). 

Jfypochlorous  Acid  as  an  Oxidizing  Agent :  Bleaching.  — 
Both  iodine  and  bromine  are  oxidized  by  hypochlorous  acid  (either  in 
pure  solution  or  in  the  form  of  chlorine  water,  equation  1,  p.  115),  the 
former  much  more  rapidly  than  the  latter,  2HOC1  +  I2  — >  2HOI  + 
C12.  Further  oxidation  to  HIO3  occurs  immediately  (see  p.  198). 
Although  iodine  has  less  affinity  for  hydrogen  than  has  chlorine 
(p.  172),  this  action  shows  that  the  relation  towards  oxygen  is  just 
the  opposite.  Here  the  iodine  goes  into  combination  and  the 
chlorine  is  displaced. 

It  is  on  account  of  its  oxidizing  power  that  hypochlorous  acid  is 
used  commercially  in  bleaching.  It  is  not  applied  to  paints,  which 
are  chiefly  mineral  substances,  but  to  complex  compounds  of  carbon, 
such  as  constitute  the  coloring  matters  of  plants  and  of  those  artifi- 
cial dyes  whose  manufacture  has  now  become  so  gigantic  an  industry. 
It  should  be  understood  that  the  great  majority  of  the  complex  com- 
pounds of  carbon  are  colorless.  Even  a  slight  chemical  change, 
affecting  only  one  or  two  of  the  atoms  in  a  complex  molecule,  is  thus 
almost  sure  to  give  a  colorless  or  much  less  strongly  colored  material. 
Indigo  (C16H10N2O2),  which  has  a  deep-blue  color,  is  an  example  of  a 
vegetable  dye  which  is  also  made  artificially.  Hypochlorous  acid 
oxidizes  it  to  isatin,  a  yellow  substance  relatively  pale  in  color: 

C16H10N2O2  +  2HOC1  ->  2C8H5N02  +  2HC1. 

In  ways  just  as  definite  as  this,  hypochlorous  acid  will  change  the 
composition  of  other  colored  substances,  although,  since  we  do  not 
know  the  formulae  of  all  these  substances,  we  cannot  always  write 
equations  for  the  actions. 

On  account  of  the  hypochlorous  acid  which  is  already  present  in 
chlorine  water,  this  solution  is  a  very  efficient  bleaching  agent.  The 
removal  of  this  one  of  the  factors  in  the  reverse  action  (p.  183) 


OXIDES   AND   OXYGEN   ACIDS   OF   CHLORINE  193 

enables  more  of  the  acids  to  be  produced  from  the  chlorine  and  water 
until  the  whole  of  the  halogen  has  been  consumed. 

As  a  rule,  bleaching  is  actually  carried  out  by  liberating 
hypochlorous  acid  from  bleaching  powder  by  means  of  sulphuric 
acid: 

OC1     H  HOC11 

Ca  '        +      )  (S04)  fc?  CaS04  +          U+  H20  +  C12. 


X 


C1  HC1 


Of  course,  temporarily,  most  of  the  hypochlorous  acid  interacts  with 
the  hydrochloric  acid  to  give  chlorine  and  water,  but,  as  the  residual 
hypochlorous  acid  loses  its  oxygen,  the  secondary  action  is  again 
displaced  backwards  until  the  chlorine  is  all  used  up. 

The  yarn  or  cloth  is  first  cleansed  from  fatty  or  oily  material  by 
boiling  with  soap  solution.  It  is  then  immersed  in  bleaching  powder 
solution,  and  finally  in  dilute  sulphuric  acid.  Both  solutions  must 
be  very  weak  in  order  that  no  interaction  may  occur  with  the  fabric 
itself.  The  last  two  processes  may  be  repeated,  if  the  brownish  or 
yellowish  coloring  material  has  not  disappeared  after  the  first 
treatment. 

Hypochlorous  acid  can  be  used  to  bleach  linen  or  cotton,  because 
the  body  of  these  materials,  apart  from  the  small  amount  of  coloring 
matter,  is  composed  of  compounds  containing  nothing  but  carbon, 
hydrogen,  and  oxygen.  These  compounds  are  very  slowly  affected 
by  hypochlorous  acid,  unless  too  strong  a  solution  is  used,  or  the 
exposure  to  its  influence  is  too  long.  That  chemical  action  does 
occur  is  shown  by  the  "  rotting  "  of  goods  which  have  not  been 
washed  thoroughly  after  bleaching.  Wool,  silk,  and  feathers,  on 
the  other  hand,  are  composed  largely  of  compounds  containing  nitro- 
gen in  addition  to  the  above  three  elements  (see  Dyeing).  Their 
constituent  material  interacts  as  easily  with  hypochlorous  acid  as 
do  the  traces  of  coloring  substances.  Hence,  since  the  fabric  itself 
would  be  attacked  by  this  agent,  different  means  of  bleaching  have 
to  be  used  for  materials  of  this  class. 

It  should  be  understood  that  a  cold  dilute  solution  of  hypochlorous 
acid  may  be  kept  almost  indefinitely  and  will  not  give  up  its  oxygen 
spontaneously.  The  transfer  takes  place  when,  and  only  when,  the 
acid  comes  in  contact  with  some  substance  capable  of  uniting  with 
oxygen. 


194  COLLEGE   CHEMISTRY 

Explanation  of  the  Activity  of  Hypo  chlorous  Addas  an  Oxi- 
dizing Agent.  —  When  hypochlorous  acid  decomposes  into  hydro- 
chloric acid  and  oxygen,  much  heat  is  liberated  (equation  1,  below). 
The  acid,  therefore,  possesses  much  more  internal  energy  than  do 
hydrogen  chloride  and  free  oxygen.  On  this  account  it  brings  to  the 
task  of  oxidizing  any  substance  more  energy  than  does  oxygen  itself, 
and  is  therefore  more  efficient.  Thus,  free  oxygen  does  not  interact 
in  the  cold  with  indigo,  or  with  bromine  or  iodine  (see  above),  while 
hypochlorous  acid  oxidizes  them*  rapidly.  The  heats  of  reaction 
show  the  difference  very  clearly.  In  equation  (2),  1800  cal.  is  the 
amount  of  heat  which  would  be  liberated  if  indigo  could  be  oxidized 
to  isatin  by  oxygen  gas.  When  hypochlorous  acid  is  used,  we 
obtain,  in  addition,  the  heat  of  decomposition  of  this  substance 
(equation  1),  so  that  the  total  heat  liberated  (equation  3),  20,400 
cal.,  is  ten  times  as  great  as  in  equation  (2)  where  free  oxygen  is  the 
oxidizing  agent : 

2HOC1  =  2HC1  (+  2O)  +  18,600  cal.  (1) 

C16H10N2O2  +  (2O)  =  2C8H5NO2  +  1800  cal. (2) 

C16H10N2O2  +  2HOC1  -  2C8H5NO2  +  2HC1  +  20,400  cal.        (3) 

By  similar  reasoning  we  explain  the  superiority  of  potassium  per- 
manganate over  free  oxygen  for  oxidizing  hydrochloric  acid  (p.  111). 
Substances  which  are  more  active  oxidizers  than  is  free  oxygen 
may  be  called  active  oxidizing  agents. 

Simultaneous,  Independent  Chemical  Changes  in  the  Same 
Substance.  —  As  we  have  seen,  hypochlorous  acid  undergoes  three 
different  changes.  Some  molecules  decompose  into  water  and  chlo- 
rine monoxide  (p.  191),  while  others  give  chloric  acid  and  hydrogen 
chloride,  and  still  others  hydrogen  chloride  and  oxygen.  Since  the 
same  molecule  cannot  undergo  more  than  one  of  these  different 
changes,  it  follows  that  the  actions  are  independent  of  one  another. 
This  is  shown  by  the  fact  that  in  sunlight  the  third  predominates, 
while  in  the  dark  it  falls  far  behind  the  second.  Since  the  relative 
quantities  of  the  products  vary,  the  several  simultaneous  actions 
cannot  be  put  in  the  same  equation.  The  fundamental  property  of 
an  equation  is  to  show  the  constant  proportions  by  weight  between 
every  pair  of  substances  in  it.  Hence  three  separate  equations  are 
required  in  the  present,  and  in  all  similar  cases  where  all  the  propor- 
tions are  not  constant  (see  Perchlorates,  p.  196).  Successive  actions, 


OXIDES  AND  OXYGEN   ACIDS  195 

like  (1)  followed  by  (2)  in  the  next  section  (cf.  p.  189),  however,  may 
be  combined  in  one  equation,  since  in  them  all  the  proportions  must 
necessarily  be  constant.  These  equations  are  interlocked,  for  (2) 
consumes  what  (1)  produces. 

Chlorates.  —  Like  hypochlorous  acid  itself,  the  hypochlorites 
turn  into  chlorates.  Thus,  when  chlorine  is  passed  into  a  warm, 
concentrated  solution  of  potassium  hydroxide,  and  particularly  when 
an  excess  of  chlorine  is  used,  the  potassium  hypochlorite  changes  into 
potassium  chlorate  KC1O3  as  fast  as  it  forms.  Since  this  action 
(equation  2)  requires  3KC1O,  the  equation  formerly  given  (p.  489) 
must  be  tripled: 

3C12  +  6KOH  ->  3KC1    +  3KC10  +  3H20.  (1) 

3KC1O->2KC1    +KC1O3  (2) 

Adding:        3C12  +  6KOH  ->  KC1O3  +  5KC1  +  3H2O. 

When  the  solution  is  cooled,  the  chlorate  crystallizes  out. 

This  action  involves  converting  five-sixths  of  the  valuable  potas- 
sium hydroxide  into  the  relatively  less  valuable  potassium  chloride. 
Hence,  in  practice,  the  makers  carry  out  the  corresponding  action 
with  calcium  hydroxide.  They  then  add  potassium  chloride  to  the 
resulting  solution,  containing  calcium  chloride  and  calcium  chlorate 
Ca(ClO3)2.  The  potassium  chlorate,  formed  by  double  decomposi- 
tion, crystallizes  when  the  solution  is  cooled. 

All  chlorates  are  at  least  moderately  soluble  in  water  (see  Table 
inside  of  front  cover).  Potassium  chlorate  is  used  in  making  fire- 
works, explosives,  and  matches.  An  intimate  mixture  with  sugar 
(C12H22011)  burns  with  semi-explosive  violence,  the  oxygen  of  the 
salt  combining  with  the  carbon  and  hydrogen  of  the  sugar  to  form 
carbon  dioxide  and  water.  Detonating  fuses  for  artillery  are  made 
of  a  mixture  of  this  salt  with  antimony  trisulphide  (q.v.). 

Chloric  Acid  HClOy  —  This  acid  may  be  obtained  in  solution 
in  water,  by  adding  the  calculated  amount  of  diluted  sulphuric  acid 
to  a  solution  of  barium  chlorate: 

Ba(ClO3)2  +  H2S04 1;  BaSO,  \  +  2HC1O3. 

The  barium  sulphate,  being  insoluble,  is  removed  by  filtration  (cf. 
p.  188). 

The  solution  may  be  concentrated  (to  about  40  per  cent)  by  evapo- 


196  COLLEGE    CHEMISTRY 

ration,  but  must  not  be  headed  above  40°,  as  the  acid  decomposes 
near  this  temperature.  The  resulting  thick,  colorless  liquid  has 
powerful  oxidizing  qualities,  setting  fire  to  paper  (made  of  cellulose 
^C8H10O5)n)  which  has  been  dipped  into  it.  It  converts  iodine  into 
iodic  acid,  2HC103  +  I2  -*  2HIO3  +  C12.  When  warmed  beyond 
40°  the  acid  decomposes,  giving  chlorine  dioxide  and  perchloric  acid: 

3HC103  ->*H2O  +  2CfO2  +  HC1O4. 

Chlorine  Dioxide  :  Chfoyous  Acid. —  Chlorine  dioxide  C102  (see 
above)  is  a  yellow  gas  which  may  be  liquefied,  and  boils  at  +  10°. 
The  gas  and  liquid  are  violently  explosive,  the  substance  being 
resolved  into  its  elements  with  liberation  of  much  heat.  It  is  formed 
whenever  chloric  acid  is  set  free,  and  hence  it  is  seen  when  a  little 
powdered  potassium  chlorate  is  touched  with  a  drop  of  concentrated 
sulphuric  acid  (end  of  last  section).  Concentrated  hydrochloric  acid 
turns  yellow  from  the  same  cause  when  any  chlorate  is  added  to  it. 
These  actions  are  used  as  tests  for  chlorates,  and  distinguish  them 
from  perchlorates  (q.v.).  With  water,  chlorine  dioxide  gives  a 
mixture  of  chlorous  acid  HC102  and  chloric  acid,  and  with  bases  a 
mixture  of  the  chlorite  and  chlorate. 

Perchlorates,  Perchloric  Acid,  and  Perchloric  Anhydride. 

-When  heated  chlorates  give  perchlorates.     Chlorates   also  give 
oxygen  at  the  same  time  (p.  47) : 

(2KC1O3->2KC1  +  302, 
(4KC1O3->3KC1O4  +  KC1. 

These  actions,  like  the  three  decompositions  of  hypochlorous  acid 
(p.  191),  are  independent,  and  proceed  simultaneously  (p.  194). 
Their  relative  speed,  however,  varies  with  the  temperature,  and  the 
decomposition  into  chloride  and  oxygen  may  completely  outrun  the 
other  when  a  catalytic  agent  like  manganese  dioxide  is  added  (p.  47). 
When  pure  potassium  chlorate  is  heated  cautiously,  about  one-fifth 
of  it  has  lost  all  its  oxygen  by  the  time  the  rest  has  turned  into  per- 
chlorate.  The  mixture  may  be  separated  by  grinding  with  the 
minimum  quantity  of  water  which  will  dissolve  the  chloride  it  con- 
tains. The  perchlorate,  having  at  15°  less  than  one-twentieth  of  the 
solubility  of  the  chloride,  will  remain,  for  the  most  part,  undissolved. 
The  perchlorates  are  much  more  stable  (p.  81)  than  the  chlorates, 


OXYGEN   ACIDS   OF   BROMINE  197 

or  hypochlorites  :  they  are  all  soluble  ii>  water,  and  they  are  used  in 
making  matches  and  fireworks. 

Pure  perchloric  acid  HC1O4  explodes  when  heated  above  92°.  But, 
like  other  liquids,  its  boiling-point  is  lower  when  its  vapor  is  under 
reduced  pressure  (cf.  p.  79).  At  56  mm.  pressure  it  boils  at  39°,  a 
temperature  at  which  hardly  any  decomposition  is  noticeable.  Hence 
the  acid  may  be  made  by  mixing  potassium  perchlorate  and  con- 
centrated sulphuric  acid  and  distilling  the  mixture  .  cautiously  in  a 
vacuum  : 

KC104  +  H2S04^KHS04  +  HClO4t- 


To  secure  the  requisite  low  pressure,  the  ordinary  distilling  apparatus 
(Fig.  14,  p.  26)  is  ma^de  completely  air-tight,  and  is  connected  by  a 
branch  tube  with  a  water-pump. 

Perchloric  acid  is  a  colorless  liquid,  which  decomposes,  and  often 
explodes  spontaneously,  when  kept.  A  70  per  cent  solution  in  water 
is  perfectly  stable,  however.  Although  it  is  an  active  oxidizing 
agent.,  it  is  not  so  active  as  chloric  acid,  and  does  not  oxidize  hydro- 
gen chloride  in  cold  aqueous  solution.  When  liberated  by  concen- 
trated sulphuric  acid  it  does  not  at  once  give  the  yellow  chlorine 
dioxide  (p.  196). 

Perchloric  anhydride  C1207  may  be  prepared  by  adding  phosphoric 
anhydride  to  perchloric  acid  in  a  vessel  immersed  in  a  freezing  mix- 
ture, P2O5  +  2HC1O4  -»  2HPO3  +  C1,O7.  Phosphoric  anhydride  is 
often  used  in  this  way  for  removing  the  elements  of  water  from  com- 
pounds. By  gently  warming  the  mixture,  the  perchloric  anhydride 
can  be  distilled  off.  It  is  a  colorless  liquid  boiling  at  82°  (760  mm.) 
and  exploding  when  struck  or  too  strongly  heated. 

Oocygen  Adds  of  Bromine.  —  No  oxides  of  bromine  have  been 
made,  but  the  acids  HBrO  (hypobromous  acid)  and  HBrO3  (bromic 
acid)  and  their  salts  are  familiar. 

By  the  action  of  bromine  on  dilute,  cold  potassium  hydroxide 
solution,  potassium  bromide^nd  hypobromite  are  formed: 

Br2  +  2KOH  '  ->  KBr  +  KBrO  +  H20. 

When  the  solution  is  heated,  the  hypobromite  turns  into  potassium 
bromate  and  bromide.  The  actions  are  exact  parallels  of  the  corre- 
sponding ones  for  chlorine  (pp.  189.  195). 

Aqueous  bromic  acid  HBrO3  may  be  made  in  the  same  way  as 


198  COLLEGE    CHEMISTRY 

chloric  acid  (p.  195),  or  by  the  action  of  chlorine  and  water  (equation 
1,  p.  115)  on  bromine: 

5HC10  +  H20  +  Br2  -» 2HBr03  +  5HC1, 

The  solution  is  colorless  and  has  powerful  oxidizing  properties. 
Thus,  it  converts  iodine  into  iodic  acid :  2HBrO3  -f  I2  — >  2HIO3  +  Br2. 
It  appears,  therefore,  that  iodine  has  more  affinity  for  oxygen  than 
has  bromine. 

• 

The  Oxide  and  Oxygen  Acids  of  Iodine.  —  The  following  are 
the  familiar  acids  and  their  corresponding  salts: 

HIO3  Iodic  acid,  KIO3         Potassium  iodate, 

[HIO4  Periodic  acid],  NaI04       Sodium  periodate, 

H5IO6  Periodic  acid,  Na2H3IO8  Disodium  periodate. 

There  is  one  oxide,  iodic  anhydride  I205. 

The  potassium  and  sodium  iodates,  KI03  and  NaI03,  are  found  in 
Chili  saltpeter.  They  may  be  made,  in  much  the  same  fashion  as 
are  the  chlorates  and  bromates  (pp.  195,  197),  by  adding  powdered 
iodine  to  a  hot  solution  of  potassium  or  sodium  hydroxide. 

Iodic  Acid  HI03  is  formed  by  passing  chlorine  through  iodine  sus- 
pended in  water.  The  action  is  parallel  to  that  of  chlorine  on  bromine 
water: 

5HC1O  +  H2O  +  I2  ->  2HIO3  +  5HC1. 

A  still  better  way  is  to  boil  iodine  with  aqueous  nitric  acid  (q.v.). 
The  latter  gives  up  oxygen  readily,  and  is  here  used  solely  on  this 
account.  Hence  it  may  be  omitted  from  the  equation,  only  the 
oxygen,  of  which  it  is  the  source,  appearing: 

I2  +  H2O  +  5O  ->  2HI03. 

In  both  these  actions  the  initial  substances  (including  the  excess  of 
nitric  acid)  and  the  products,  with  the  exception  of  the  iodic  acid 
itself,  are  all  volatile.  When  the  solution  is  concentrated  by  evapora- 
tion, therefore,  only  the  iodic  acid  crystallizes.  It  is  a  white  solid, 
perfectly  stable  at  ordinary  temperatures,  andean  be  kept  indefinitely. 
At  170°  it  begins  to  give  off  water  vapor  (2HI03  <=>  H2O  +  I2O5), 
leaving  iodic  anhydride.  The  latter  is  a  white  crystalline  powder 
which  may  be  raised  to  300°  before  it,  in  turn,  breaks  up,  giving 
iodine  and  oxygen. 

Some  sodium  periodate  (NalOJ  is  found  in  Chili  saltpeter. 


OXIDES   AND   OXYGEN   ACIDS,   THE    HALOGENS  199 

Chemical  Relations.  —  The  compounds  of  the  halogens  with 
metals  and  with  hydrogen  diminish  in  stability,  with  ascending 
atomic  weight  of  the  halogen,  in  the  order:  F(19),  Cl(35.5),  Br(80), 
I  (127).  Each  halogen  will  displace  those  following  it  from  this 
kind  of  combination.  In  the  case  of  the  oxygen  compounds,  the 
order  of  stability  is  just  the  reverse,  those  of  iodine,  for  example, 
being  the  only  ones  which  are  reasonably  stable.  The  order  of 
displacement  in  such  compounds  confirms  this  conclusion. 

Amongst  the  oxygen  acids  of  any  one  halogen,  those  containing 
most  oxygen  are  most  stable.  The  salts  are  in  all  cases  more  stable 
by  far  than  the  corresponding  acids. 

The  halogens  when  combined  with  metals  and  hydrogen  are  univa- 
lent  (HI,  KC1,  etc.).  It  is  clear,  however,  that,  when  united  with 
oxygen,  their  valence  is  higher.  The  maximum  is  shown  in  per- 
chloric anhydride  (C1207),  where  chlorine  appears  to  be  heptavalent. 

The  formulas  of  the  acids  might  be  written  so  as  to  retain  the 
univalence : 

H— Cl,  H-O-C1,  H-O-0-C1,  H-O-0-O-C1, 
H-O-O-O-0-C1. 

But  compounds  in  which  we  are  compelled  to  believe  that  two  oxygen 
units  are  united  are  usually  unstable  (see  Hydrogen  peroxide),  and 
we  should  expect  the  instability  would  be  greater  with  three  and 
with  four  units  of  oxygen  in  combination.  Here,  however,  the 
reverse  state  of  affairs  must  be  taken  account  of  in  our  formulae,  for 
HC104  is  the  most  stable  of  the  chlorine  set.  This  reasoning, 
together  with  the  heptavalence  in  C12O7,  leads  us  to  assume  the 
valence  seven  in  perchloric  acid  (see  Periodic  system).  The  struc- 
tural formulae  (cf.  p.  156)  of  some  of  these  substances  are  therefore 
often  written  as  follows: 

O  O 

II  II 

H-C1,     H-O-C1,     H-0-C1-O,     Na-O-I  =  0. 

II  II 

O  O 

Exercises.  —  1.  Assign  to  its  proper  class  (pp.  124,  163,  189) 
each  of  the  actions  mentioned  in  this  chapter. 

2.  Knowing  that  potassium  fluosilicate  (K2SiF6)  is  insoluble,  how 
should  you  make  chloric  acid  (p.  188)? 

3.  Make  the  equation  for  the  interaction  of  chlorine  with  calcium 


200  COLLEGE   CHEMISTRY 

hydroxide  in  hot  water  (p.  195).    How  should  you  make  zinc  chlorate 
from  zinc  hydroxide  (Zn(OH)2)? 

4.  How  should  you  make  pure  potassium  hypochlorite  from 
hypochlorous  acid  (p.  191)? 

5.  On  what  circumstances  would  the  possibility  of  making  barium 
chlorate  by  action  of  chlorine  on  barium  hydroxide  depend  (p.  195)? 
Could  pure  barium   chlorate   be  obtained  easily  by   this   means 
(see  Table  of  solubilities)? 

6.  Make  the  equations  for:  (a)  the  preparation  of  potassium  bro- 
mate;  (6)  pure  aqueous  bromic  acid;  (c)  the  interaction  of  iodine 
with  aqueous  potassium  hydroxide  in  the  cold,  and  when  heated. 

7.  Make  the  equations  for  the  interactions  of  chlorine  dioxide  with 
water,  and  with  aqueous  potassium  hydroxide. 


, 


CHAPTER  XVII 
DISSOCIATION    IN    SOLUTION 

THE  employment  of  interacting  substances  in  the  form  of  solutions 
is  so  constant  in  chemistry,  and  the  reasons  for  this  are  so  cogent, 
that  we  must  now  resume  the  discussion  of  the  subject  of  solution 
(cf.  p.  96). 

The  present  chapter  will  be  devoted  to  giving  the  proofs  that,  to 
speak  in  terms  of  the  molecular  hypothesis,  the  molecules  of  acids, 
bases,  and  salts  in  aqueous  solutions,  are  actually  dissociated  into  parts 
by  the  solvent.  This  will  be  shown  by  consideration,  successively, 
of  certain  peculiarities  in  the  chemical  behavior,  and  the  freezing- 
points  of  the  solutions  of  these  substances.  We  shall  see  that  these 
parts  coincide  in  composition  with  the  radicals. 

Some  Characteristic  Properties  of  Acids,  Bases,  and  Salts, 
Shown  in  Aqueous  Solution.  —  Acids  all  contain  hydrogen  (p.  64). 
In  aqueous  solution,  if  soluble,  they  are  so  T  in  taste,  they  turn  blue 
litmus  red,  and  their  hydrogen  is  displaced  by  certain  metals  (p. 
65),  and  has  the  properties  of  a  radical.  By  the  last  statement  is 
meant  that  it  very  readily  exchanges  places  with  other  radicals  in 
reversible  double  decompositions  (p.  187).  Many  other  bodies,  like 
sugar,  kerosene,  and  alcohol,  contain  hydrogen  also,  but  not  one.  of 
them  shows  all  of  these  properties.  Again,  all  salts  are  made  up  of 
two  radicals,  and  the  reversible  double  decompositions  into  which 
they  enter  with  acids,  bases,  and  other  salts,  consist  in  exchanges  of 
these  radicals.  Other  substances  may  include  the  same  combinations 
of  atoms,  but  in  their  actions  these  groupings  are  often  disregarded. 
Thus,  sodium  chloride  and  silver  nitrate  exchange  radicals  com- 
pletely (p.  187),  and,  in  dilute  solution,  hydrogen  chloride  and  sodium 
hydrogen  sulphate  do  so  partially  (p.  118).  But  sodium  chloride 
and  nitroglycerine  C3H6(N03)3  do  not  interact  at  all.  The  latter  is 
not  a  salt,  although  it  contains  the  same  proportion  of  nitrogen  to 
oxygen  as  does  any  nitrate. 

Furthermore,  it  is  chiefly  in  aqueous  solution  that  these  special  prop- 

201 


202  COLLEGE   CHEMISTRY 

erties  of  acids,  bases,  and  salts  become  apparent.  Their  behavior 
is  often  quite  different  in  the  absence  of  this  solvent.  If,  for  example, 
we  mix  together  ammonium  carbonate  and  partially  dehydrated 
cupric  nitrate,  and  apply  heat,  a  violent  interaction  begins.  An 
immense  cloud  of  smoke  and  gas  is  thrown  out  of  the  tube,  and  the 
substance  remaining  is  either  black,  or  reddish,  in  parts,  according  to 
the  proportions  of  the  substances  employed.  The  residue  contains 
cupric  oxide,  and  sometimes  red  cuprous  oxide  (Cu20).  The  gas  is 
tinged  red  by  the  presence  of  nitrogen  tetroxide  (NO2),  while  a  more 
careful  examination  would  show  that  it  contained  carbon  dioxide, 
nitrogen,  nitrous  oxide  (N2O),  water  vapor,  and  perhaps  still  other 
products.  The  contrast,  when  the  substances  are  dissolved  in  water 
before  being  brought  in  contact  with  one  another,  is  very  great.  A 
pale-green  precipitate  is  formed  at  once,  and  rapidly  settles  out. 
On  examination,  this  turns  out  to  be  a  carbonate  of  copper  (basic), 
while  evaporation  of  the  solution  furnishes  us  with  ammonium 
nitrate.  There  are  only  two  main  products,  and  the  essential 
part  of  the  action  in  solution  may  be  represented  by  the  equation  : 

(NH4)2C03  +  Cu(N03)2^CuC03  +  2NH4N03. 

In  the  interaction  between  the  dry  substances  the  molecules  are  com- 
pletely disintegrated,  and  the  whole  change  is  very  complex.  In  the 
action  in  water  no  heating  is  required,  the  substances  are  neatly 
broken  apart,  certain  groups  of  atoms,  which  we  call  radicals,  are 
transferred  as  wholes  from  one  state  of  combination  to  another,  and 
the  rearrangement  takes  place  in  a  machine-like  manner.  Contrasts 
like  this  between  the  interactions  of  anhydrous  and  dissolved  bodies 
are  very  common. 

Many  compounds,  however,  do  not  show  any  change  in  behavior 
when  dissolved  in  water.  Sugar,  for  example,  is,  as  a  rule,  more 
readily  acted  upon  in  the  absence  of  any  solvent.  Then  again,  while 
water  is  not  the  only  solvent  which  has  the  effect  we  have  just 
described,  the  majority  of  solvents,  if  they  affect  chemical  change  at 
all,  simply  retard  it.  Thus  the  union  of  iodine  and  phosphorus  in 
the  absence  of  a  solvent  takes  place  spontaneously  with  a  violent 
evolution  of  heat.  When  the  elements  are  dissolved  in  carbon  bisul- 
phide, before  being  mixed,  the  action  is  much  milder,  although  the 
product  is  the  same  (phosphorus  tri-iodide).  The  diminution  in  the 
concentration  of  the  ingredients  has  decreased  the  speed  of  the  action 


DISSOCIATION   IN  SOLUTION  203 

in  the  normal  way  (p.  180).  That  water  and  some  other  solvents 
have  a  specific  influence  tending  to  increase  the  apparent  activity  of 
certain  classes  of  substances,  shows  that  a  special  explanation  of  the 
phenomenon  must  be  found. 

Summing  up  these  points  we  see  that  the  peculiarity  of  acids, 
bases,  and  salts  in  aqueous  solution  is  that  each  compound  always 
splits  in  the  same  way.  Thus,  cupric  nitrate  always  gives  changes 
involving  Cu  and  NO3  and  never  interacts  so  as  to  use  CuN2  and  O3, 
or  CuO2  and  NO2,  as  the  basis  of  exchange.  Similarly,  dilute  acids 
always  offer  hydrogen  in  exchange,  and  so  nitric  acid  behaves  as  if 
composed  of  H  and  N03,  and  sulphuric  acid  as  if  composed  of  2H 
and  SO 4,  and  never  as  if  made  up  of  HSO  and  HO3,  or  H2S  and  O4. 
The  sour  taste  and  the  effect  upon  litmus  seem  to  be  properties  of 
this  easily  separable  hydrogen,  for  they  are  shown  only  by  acids. 
The  result  is  that  we  can  make  a  list  of  the  units  of  exchange,  such 
as  H,  OH,  NO3,  CO3,  SO4,  Cu,  K,  and  Cl,  employed  by  acids,  bases, 
and  salts  in  their  interactions.  The  molecule  of  each  compound  of 
these  classes  contains  at  least  two  of  them.  Even  when  these  units 
contain  more  than  one  atom,  their  coherence  is  as  noticeable  within 
this  class  of  actions,  as  is  the  permanence  of  the  atomic  masses 
themselves  in  all  actions. 

The  question  raised  in  our  minds  is  whether  solution  in  water 
alters  the  character  of  the  molecule,  simply  by  producing  a  sort  of 
plane  of  cleavage  in  it  which  creates  a  predisposition  to  a  uniform 
kind  of  chemical  change,  or  whether  it  actually  divides  the  molecules 
into  separate  parts  consisting  of  the  above  units  of  exchange,  and 
leaves  subsequent  chemical  actions  to  occur  by  cross-combination  of 
these  fragments.  The  fact  that  the  dissolved  substances  can  be 
recovered  by  evaporation  of  the  liquid  does  not  demonstrate  that 
they  have  not  been  decomposed  temporarily  while  in  solution.  The 
alteration  which  the  water  produces,  whatever  it  be,  will  naturally 
be  reversed  when  the  water  is  removed.  Since  our  question  involves 
nothing  but  the  counting  of  particles,  the  number  oT  which  would 
be  much  greater  in  the  event  that  actual  subdivision  of  molecules 
is  the  explanation,  it  can  be  answered  by  a  study  of  the  physical 
properties  of  solutions.  Several  physical  properties  can  be  used, 
and  they  give  concordant  answers  to  the  question.  We  shall  con- 
fine ourselves  here,  however,  to  the  evidence  furnished  by  the 
freezing-points  of  solutions. 


204  COLLEGE   CHEMISTRY 

Freezing -Points  of  Solutions.  — Every  pure  liquid  has  a  def 
nite  temperature  at  which  it  freezes.  Thus,  pure  water  freezes  a 
0°  and  benzene  at  6°.  The  presence  of  a  foreign,  dissolved  bod] 
however,  lowers  the  freezing-point.  Thus,  sea-water  is  harder  t 
freeze  than  fresh  water.  The  freezing-points  of  solutions  of  the  sam 
substance  are  found  to  be  depressed  below  that  of  the  solvent  in  prc 
portion  to  the  concentration  of  the  solute.*  Thus,  in  one  set  c 
experiments,  solutions  of  sugar  containing  11.4  g.,  22.8  g.,  and  34. 
g.  of  sugar  to  100  g.  of  water  were  found  to  freeze  at  —  0.62°,  —  1.23' 
and  —  1.85°,  respectively. 

In  everyday  life  we  scatter  salt  on  ice  to  melt  it.  The  salt  du 
solves  in  the  moisture  on  the  surface,  and  the  ice  cannot  exist  i 
presence  of  this  solution  at  0°  (p.  78).  It  melts,  absorbing  heat  i 
doing  so,  until  the  temperature  of  the  mixture  reaches  that  of 
freezing,  saturated  solution  of  salt  (about  —  21°  =  —  6°  F.).  Whe 
the  existing  temperature  is  lower  than  this,  the  salt  has  no  effect  o 
the  ice.  Freezing  mixtures,  being  usually  mixtures  of  ice  with  variou 
soluble  substances,  work  in  accordance  with  this  principle. 

Laivs  of  Freezing-Point  Depression.  —  The  depression  in  th 
freezing-point  is  directly  proportional  to  the  weight  of  dissolve 
substance  in  a  given  amount  of  the  solvent.  The  depression  : 
inversely  proportional  to  the  amount  of  solvent.  Thus,  if  we  doubl 
the  concentration  of  the  solution,  the  depression  in  the  freezing 
point  is  doubled.  Further,  equal  numbers  of  molecules  of  differei 
solutes  in  the  same  quantity  of  solvent  give  equal  depressions.  Or,  i 
other  words,  the  depression  is  proportional  to  the  concentration  c 
the  molecules  of  the  solute.  Thus,  solutions  containing  342  g.  c 
sugar  (C12H22Out),  or  46  g.  of  alcohol  (C2H60),  or  74  g.  of  metrrj 
acetate  CH3(C2H3O2),  in  1000  g.  of  water,  being  weights  which  cor 
tain  equal  numbers  of  molecules,  show  a  depression  below  the  free; 
ing-point  of  water  of  about  1.89°  in  each  case.  That  is,  such  solutior 
all  freeze  close  to  —  1.89°.  This  depression,  produced  by  a  mole  c 
the  solute  in  1  1.  of  water,  is  called  the  molecular  depression  constan 
and  has  a  different  value  for  each  solvent.  For  solutions  of  the  sam 

*  The  ice  which  separates  during  the  freezing  consists,  as  a  rule,  of  the  pui 
solvent,  and  the  solute  does  not  enter  into  it.  Only  when  this  is  the  case  do< 
this  law  represent  the  facts. 

t  12  x  12  +  22  x  1  +  11  X  16  -  342. 


i 


O  v\. 


1.89°  x 


Wt.  of  Solute  1000 


Mol.  Wt.  of  Solute       Wt.  of  Solvent. 


•'    '     / 

DISSOCIATION   IN   SOLUTION  205 

t~0         *?** 

molecular  concentration  in  benzene  the  depression  isj.9^  in  phenol 
(carbolic  acid)  7.5°.     Combining  these  facts  in  one  expression: 
The  observed'  depression 
in  an  aqueous  solution 

/  '"»  <   £  £) 

For  other  solvents,  the  corresponding  value  of  the  depression  con- 
stant must  be  substituted  for  1.89°. 

These  laws  describe  the  facts  most  exactly  when  the  solutions  are 
dilute.  They  hold  only  when  there  is  no  chemical  interaction 
between  solute  and  solvent.  Even  so,  however,  acids,  bases,  and 
salts  dissolved  in  water  present  many  apparent  exceptions  and  must  be 
discussed  separately  (see  below). 

It  will  be  noted  that,  when  the  other  factors  in  the  foregoing  equa- 
tion are  known  or  observed,  the  molecular  weight  of  the  solute  may  be 
determined.  The  fact  makes  possible  the  determination  of  this 
constant  for  substances  which  are  not  volatile  (see  Hydrogen 
peroxide). 

Freezing- Points  and  Dissociation  in  Solution.  —  The  sub- 
stances which  present  the  most  conspicuous  exceptions  to  the  above 
rules  are  acids,  bases,  and  salts  in  aqueous  solution.  With  most  of 
these,  the  depression  produced  is  greater  than  we  should  expect 
from  the  concentration  of  the  solution.  Thus,  in  an  actual  experi- 
ment, two  equi-molar  solutions  were  compared.  One  contained  one 
mole  (74  g.)  of  methyl  acetate,  and  the  other  one  mole  (58.5  g.)  of 
sodium  chloride,  each  dissolved  in  2000  g.  (2  liters)  of  water.  The 
freezing-points  observed  were: 

Pure  water 0.000°       Pure  water 0.000° 

Sol.  of  methyl  acetate     .    -0.970°       Solution  of  salt     .    .    .    -1.678° 

Depression 0.970°       Depression 1.678° 

0.970° 


Excess  depression  by  salt  0 . 708° 

The  solution  of  methyl  acetate,  as  it  contained  only  0.5  moles  of 
the  solute  per  liter  of  water,  showed,  as  it  should  do,  about  half  the 
average  molecular  depression  (1.89°,  p.  204).  This  is  typical  of  the 
class  of  substances  showing  normal  behavior.  Sugar,  alcohol,  and 
hundreds  of  other  substances,  in  solutions  of  the  same  molar  con- 
centration, would  have  given  the  same  value. 

The  freezing-point  of  the  salt  solution,  however,  was  much  lower. 


206  COLLEGE   CHEMISTRY 

If  this  solution  had  really  contained  the  same  concentration  of  c 
solved  molecules  as  the  other  solution,  its  depression  would  hi 
been  0.970°  likewise.  The  number  of  molecules  in  the  solution  m 
therefore  have  been  greater  than  we  should  have  expected  from 
number  of  molecules  taken.  In  other  words,  a  portion  of  the  me 
cules  of  the  salt  must  have  been  broken  up,  and  the  excess  depressi 
0.708°,  must  have  been  due  to  the  extra  molecules  produced  by  c 
sociation.  Now  sodium  chloride  molecules  cannot  give  more  tl 
two  particles  each,  and  the  depression  is  proportional  to  the  nurn 
of  particles.  It  follows,  therefore,  that  |f  §,  or  0.732  (73.2  per  ce 
of  the  molecules  were  dissociated. 

This  result  is  typical  also.  Acids,  bases,  and  salts,  of  which  ( 
mole  is  dissolved  in  two  liters  of  water,  are  found  to  give  irregu 
values,  all  more  or  less  in  excess  of  0.970°.  Those  which  contain  t 
two  radicals,  like  sodium  chloride  (NaCl)  and  potassium  nitr; 
(KN03),  give  values  between  0.970°  and  2  X  0.970°.  Substan 
like  calcium  chloride  (CaCl2)  and  sodium  sulphate  (Na.,SO4)  g 
depressions  approaching  three  times  the  normal  value:  their  me 
cules  contain  three  radicals.  The  excess  depression  depends,  the 
fore,  upon  the  number  of  particles  which  each  molecule  can  furni 
and  upon  the  proportion  of  all  the  molecules  which  is  dissociated  ii 
these  fragments. 

In  the  case  of  an  acid,  base,  or  salt,  the  depression  is  not  stric 
proportional  to  the  concentration.  Thus,  one  mole  of  salt  in  ft 
liters  of  water  does  not  give  half  the  depression  of  the  two-li 
solution  (0.839°)  but  somewhat  more  (about  0.844°).  The  sa: 
method  of  calculation  indicates,  therefore,  a  greater  degree  of  d 
sociation  (about  79  per  cent)  in  the  more  dilute  solution  (see  lo 
equilibrium) . 

Acids,  bases,  and  salts,  so  far  as  they  are  soluble  in  materials  li 
toluene,  benzene,  chloroform,  and  carbon  bisulphide,  exhibit  simj 
normal  depressions  in  these  solvents.  It  appears,  therefore,  th 
in  many  solvents,  dissociation  does  not  take  place.  In  comm 
experience  it  is  encountered  only  in  solutions  in  water,  and,  perhaj 
alcohol. 

The  freezing-point  is  only  one  of  four  properties  of  solutions  whi 
can  be  used  for  determining  the  numbers  of  molecules  presei 
Numerous  measurements  show  that  aqueous  solutions  of  acids,  bas 
and  salts  not  only  have  abnormal  freezing-points,  but  also  abnorn 


DISSOCIATION   IN   SOLUTION  207 

osmotic  pressures  (cf.  p.  101)  and  abnormal  boiling-points.  The 
electrical  conductivity  is  the  fourth  property  which  gives  the  required 
information  (see  Chap.  xix).  Now,  when  we  observe  the  behavior 
of  the  same  solution  in  each  of  these  four  ways,  and  calculate  the 
degree  of  dissociation  from  the  result  of  each  measurement,  we  find 
that  the  values  obtained  are  usually  identical,  within  the  limits  of 
error  to  which  the  methods  are  liable.  Thus  the  indications  of  dis- 
sociation found  in  the  chemical  behavior  of  acids,  bases,  and  salts 
(p.  203)  are  fully  confirmed  by  a  study  of  the  physical  properties  of 
their  solutions.* 

Applications:  The  Constitution  of  Solutions  of  Acids, 
Bases,  and  Salts.  —  The  composition  of  solutions  which  are  nor- 
mal or  abnormal,  in  respect  to  osmotic  pressure,  freezing-point,  and 
boiling-point,  may  be  shown  thus: 


Solutes. 

Dissolved  in 
Water,  Alcohol, 
etc. 

Dissolved  in 
Toluene,  Chlo- 
roform, etc. 

Acids,  bases,  salts    

Abnormal 

Normal 

Other  substances 

Normal 

Normal 

It  appears  that  water  and  some  other  solvents  have  the  power  of 
decomposing  acids,  bases,  and  salts.  Such  solvents  have,  in  fact, 
an  effect  on  these  materials  that  resembles,  outwardly  at  least,  the 
effect  which  heat  has  on  many  substances  (e.g.  p.  81),  they  cause 
dissociation: 

NaCl<=±(Na)  +  (Cl). 

In  consequence  of  this,  our  view  of  the  nature  of  an  aqueous  solution 
of  hydrogen  chloride  (HC1),  or  common  salt  (NaCl),  or  sodium 
hydroxide  (NaOH),  or  any  of  the  substances  of  the  classes  which 
these  represent,  may  now  be  stated  in  definite  terms.  Such  a  solu- 
tion contains,  besides  undivided  molecules  of  the  solute,  at  least  two 
other  kinds  of  material,  H,  Na,f  Cl,  OH,  etc.,  which  result  from  the 

*  Recent  observations,  showing  that  in  some  cases  rapid  double  decom- 
positions of  the  normal  kind  take  place  in  solutions  which  exhibit  no  physical 
evidence  of  the  existence  of  dissociation,  demonstrate  that  it  would  have  been 
unsafe  to  infer  dissociation  from  chemical  evidence  alone. 

f  The  objection  that  separate  atoms  of  sodium  could  not  remain  free  in 
water,  will  be  disposed  of  later. 


208  COLLEGE   CHEMISTRY 

breaking  up  of  the  molecules.  We  shall  see  that  these  subdivisions 
of  the  original  molecules  have  distinct  physical  and  chemical  prop- 
erties of  their  own.  The  descriptions  of  the  "  properties  "  of  the 
solutions,  as  they  used  to  be  given  in  chemistry,  were  really  a  con- 
fused statement  of  the  properties  of  the  different  components  of  a 
mixture. 

The  free  radicals,  of  whose  existence  we  have  thus  become  con- 
vinced, constitute  a  new  set  of  materials  (with  appropriate  names. 
See  p.  224).  Thus  the  hydrogen  radical  of  acids,  although  a  form  of 
uncombined  hydrogen,  differs  totally  from  the  gas  which  is  com- 
posed of  the  same  material.  The  gas  has  no  sour  taste  or  effect 
upon  litmus:  these  are  properties  of  the  free  radical.  The  gas  is 
very  slightly  soluble  in  water,  while  the  hydrogen  radical  exists  as  a 
separate  substance  only  in  solution.  Again,  substances  with  the 
composition  of  the  radicals  NO3  and  SO4  are  not  known  at  all 
except  in  solutions. 

Exercises. —  1.  What  depression  in  the  f.-p.  of  water  will  be 
produced  by  dissolving  10  g.  of  bromine  in  1  kg.  of  this  solvent? 

2.  What  depressions  in  the  f.-p.  of  benzene  and  of  phenol  would 
be  produced  by  10  g.  of  bromine  to  1  kg.  of  the  solvent,  if  no  chemical 
action  took  place? 

3.  What  is  the  molecular  depression-constant  of  a  solvent  in 
which  5  g.  of  iodine  in  500  g.  of  the  solvent  lowers  the  f.-p.  0.7°? 

4.  What  is  the  degree  of  dissociation  of  zinc  sulphate,  if  5  g.  of  it 
dissolved  in  125  g.  of  water  produce  a  lowering  of  0.603°  in  the  f.-p.? 


•  / 


<far«" 

v/£ 


CHAPTER  XVIII 
OZONE   AND   HYDROGEN   PEROXIDE 

A  FRESH,  penetrating  odor,  resembling  that  of  very  dilute  chlorine, 
was  noticed  by  van  Marum  (1785)  as  being  perceptible  near  an  elec- 
trical machine  in  operation.  Schonbein  (1840)  showed  that  the  odor 
was  that  of  a  distinct  substance,  which  he  named  ozone  (Gk.  o£c«/, 
to  smell),  and  he  discovered  a  number  of  ways  of  obtaining  it.  It  is 
very  questionable  whether  there  is  any  ozone  in  the  air,  excepting 
temporarily  in  the  immediate  neighborhood  of  a  natural  or  artificial 
discharge  of  electricity. 

Preparation  of  Ozone  O3.  —  The  most  satisfactory  way  of 
preparing  ozone  is  to  allow  electric  waves  to  pass  through  oxygen. 
The  apparatus  (Fig.  50)  consists  of  two  co-axial  glass  tubes,  between 


FIG.  50. 


which  the  oxygen  flows.  The  waves  are  generated  by  connecting  an 
outer  layer  of  tinfoil  on  the  outer  tube,  and  an  inner  layer  of  tinfoil 
in  the  inner  tube  with  the  poles  of  an  induction  coil.  With  dry,  cold 
oxygen,  about  7.5  per  cent  of  the  gas  is  turned  into  ozone. 

Ozone  is  found  in  the  oxygen  generated  by  electrolysis  of  dilute 
sulphuric  acid  (see  p.  216).  It  arises  during  the  slow  oxidation  of 
phosphorus  by  the  air,  resulting,  probably,  from  the  decomposition 
of  unstable,  highly  oxidized  bodies  which  are  formed  during  the  action. 
Ozone  is  formed  also  when  a  jet  of  burning  hydrogen,  or  an  elec- 
trically heated  loop  of  platinum  wire  is  immersed  in  liquid  oxygen. 
Oxygen  containing  as  much  as  15  per  cent  of  it  is  produced  by  the 
interaction  of  fluorine  and  water  (p.  170). 

209 


210  COLLEGE  CHEMISTRY 

Physical  Properties  of  Ozone.  —  Ozone  is  a  gas  of  blue  color. 
It  boils  at  —  119°,  so  that  when  a  mixture  of  oxygen  and  ozone  is  led 
through  a  U-tube  immersed  in  liquid  oxygen  (—  182.5°),  the  ozone 
collects  in  the  tube  as  an  opaque,  deep-blue  fluid.  Ozone  is  much 
more  soluble  in  water  than  is  oxygen.  At  12°,  100  volumes  of  water 
would  dissolve  50  volumes  of  the  gas  at  one  atmosphere  pressure. 

Chemical  Properties  of  Ozone.  —  The  density  of  ozone  is 
one-half  greater  than  that  of  oxygen.  Its  molecular  weight  is 
therefore  48,  and  its  formula  O3.  Ozone  is  relatively  stable  when 
mixed  with  much  oxygen.  When  the  ozonized  oxygen  is  heated, 
however,  the  ozone  is  decomposed  at  about  250-300°.  The  action 
for  its  formation: 

3O2  <=±  2O3 

is  therefore  reversible.  As  the  equation  shows,  three  volumes  of 
oxygen  give  two  of  ozone.  There  is  also  an  absorption  of  much 
energy  from  the  electric  waves,  or,  in  other  methods  of  making  it, 
from  the  concomitant  chemical  changes:  O  +  O2  =  O3  —  32,400  cal. 
Ozone  is  a  much  more  active  oxidizing  agent  than  oxygen.  Mer- 
cury and  silver,  which  are  not  affected  by  the  latter,  are  converted 
into  oxides  by  the  former.  Silver  gives  the  peroxide,  Ag2O2.  Paper 
dipped  in  starch  emulsion  containing  a  little  potassium  iodide  is 
used  as  a  test  for  ozone : 

03  +  2KI  +  H20  -» 02  +  2KOH  +  I2. 

The  iodine  gives  a  deep-blue  color  to  the  starch  (cf.  p.  165).  This 
test,  however,  will  not  distinguish  ozone  from  chlorine  or  hydrogen 
peroxide,  and  may,  therefore,  be  used  only  in  the  absence  of  these 
substances.  Ozone  also  removes  the  color  from  organic  dyes,  such 
as  indigo,  by  oxidizing  them  (cf.  p.  192).  Its  activity  as  an  oxidizing 
agent,  like  the  similar  activity  of  hypochlorous  acid,  is  due  to  the 
fact  that  it  contains  much  more  energy  than  oxygen. 

Ozone  is  used  commercially  in  bleaching  oils  and  in  purifying 
starch.  It  is  employed  also  for  sterilizing  drinking  water  in  Lille 
and  other  cities. 

HYDROGEN  PEROXIDE  H2O2. 

Hydrogen  peroxide  is  found  in  minute  amounts  in  rain  and  snow. 
It  is  formed  in  small  quantities,  in  a  way  not  at  present  fully  under- 
stood, when  moist  metals  rust. 


HYDROGEN    PEROXIDE  211 

Preparation  of  Hydrogen  Peroxide. — When  sodium  peroxide 
(q.v.)  is  added,  a  little  at  a  time,  to  a  dilute  acid,  hydrogen  peroxide 
is  set  free  and  remains  dissolved  in  the  liquid. 

NaA  +  2HClt=;2NaCl  +  H202. 

When  hydrated  barium  peroxide  (Ba02,8H20)  is  shaken  with  cold, 
dilute  sulphuric  acid  a  similar  action  takes  place: 

Ba02  +  H2S04  <=?  BaSO4 1  +  H202. 

Phosphoric  acid  is  largely  employed  instead  of  sulphuric  acid  in  the 
commercial  manufacture  of  hydrogen  peroxide,  and  great  care  is 
taken  to  precipitate  the  other  products  and  all  impurities  from  the 
solution. 

An  aqueous  solution  is  also  obtained  by  passing  carbon  dioxide 
through  barium  peroxide  suspended  in  water: 

Ba02  +  C02  +  H2O  iq>  BaCO3|  +  H202. 

Pure  hydrogen  peroxide  is  isolated  from  any  of  these  solutions  by 
distillation  under  reduced  pressure  (p.  197).  It  is  much  less  volatile 
than  water,  but  decomposes  into  water  and  oxygen  violently  at  100°. 
Hence  the  lower  pressure  is  required  to  make  possible  its  volatiliza- 
tion at  a  temperature  below  this  point.  At  68  mm.  pressure,  the 
water  begins  to  pass  off  first  (at  about  45°).  The  last  portion  of  the 
liquid  boils  at  84-85°  and  is  almost  all  hydrogen  peroxide. 

By  evaporating  the  commercial  (3  per  cent)  solution  at  70°,  a 
liquid  containing  45  per  cent  of  hydrogen  peroxide  may  be  made 
without  much  loss  of  the  material  by  volatilization. 

Physical  Properties.  —  Hydrogen  peroxide  is  a  syrupy  liquid 
of  sp.  gr.  r.5.  It  blisters  the  skin,  and,  when  diluted,  has  a  disagree- 
able metallic  taste.  It  has  been  frozen  (m.-p.  —  2°). 

Chemical  Properties.  —  Hydrogen  peroxide  (100  per  cent)  is 
very  unstable,  and  decomposes  slowly  even  at  —  20°.  The  dilute 
aqueous  solution,  when  free  from  impurities,  keeps  fairly  well.  The 
presence  of  a  trace  of  free  acid  increases  its  stability.  Free  alkalies 
and  most  salts  assist  the  decomposition;  hence  the  necessity  for 
purifying  the  commercial  solution.  Addition  of  powdered  metals, 
of  manganese  dioxide,  or  of  charcoal  causes  effervescence  even  in 
dilute  solutions,  and  oxygen  escapes: 

2H2O2  ->  2H2O  +  O2. 


212  COLLEGE   CHEMISTRY 

Since  the  substance  cannot  be  vaporized,  even  at  low  pressure, 
without  some  decomposition,  its  molar  weight  has  been  determined 
by  the  freezing-point  method.  The  freezing-point  of  a  3.3  per  cent 
solution  in  water  was  —  2.03°.  Substitution  of  these  data  in  the 
formula  (p.  205)  gives  31.8  g.  as  the  molar  weight.  Now  the  formula 
HO  corresponds  to  a  molar  weight  of  17  and  H2O2  to  one  of  34.  It 
is  evident,  therefore,  that  the  latter  is  the  correct  formula. 

Hydrogen  peroxide,  in  solution  in  water,  is  a  feeble  acid.  As  an 
acid  it  enters  into  double  decomposition  readily,  and  the  peroxides 
are  really  salts  in  which  the  negative  radical  is  02".  Thus,  when 
hydrogen  peroxide  is  added  to  solutions  of  barium  and  strontium 
hydroxides,  the  hydrated  peroxides  appear  as  crystalline  precipitates : 

Sr(OH)2  +  H2O2  <±  2H20  +  Sr02. 

The  precipitation  involves  another  equilibrium:  Sr02+8H20  +±  Sr02, 
8H20  (solid). 

The  formation  of  a  beautiful  blue  substance  by  the  action  of 
hydrogen  peroxide  upon  dichromic  acid  is  used  as  a  test.  The  test 
is  carried  out  by  adding  a  drop  of  potassium  dichromate  to  an 
acidulated  solution  of  the  peroxide.  The  acid  interacts  with  the 
dichromate,  giving  free  dichromic  acid: 

H2S04  +  K2Cr2O7^±H2Cr207  +  K2S04. 

The  composition  of  the  blue  substance,  which  is  very  unstable  and 
quickly  decomposes,  is  not  certainly  known,  so  that  no  equation  for 
its  formation  can  be  given.  The  blue  substance  has  the  property, 
unusual  in  inorganic  compounds,  of  dissolving  much  more  readily 
in  ether  than  in  water.  It  is  also  much  less  unstable  when  removed 
from  the  foreign  materials  in  the  aqueous  solution.  Hence  the  test 
is  rendered  more  delicate  by  extracting  the  solution  with  a  small 
amount  of  ether.  In  the  ethereal  layer  the  color  of  the  compound  is 
more  permanent,  as  well  as  more  distinctly  visible  on  account  of  the 
greater  concentration. 

Hydrogen  peroxide  is  a  much  more  active  oxidizing  agent  than  is 
free  oxygen.  This  would  be  expected  from  the  fact,  that  it  contains 
so  much  more  internal  energy  than  the  water  and  oxygen  into  which 
it  decomposes  (p.  194),  that  23,100  cal.  are  liberated  in  the  decom- 
position of  one  mole.  Thus,  it  liberates  iodine  from  hydrogen  iodide, 
an  action  which,  in  presence  of  starch  emulsion  (cf.  p.  165),  is  used 
as  a  test  for  its  presence: 

2HI  +  H202  ->  2H20  +  I2. 


HYDROGEN  PEROXIDE  213 

It  converts  sulphides  into  sulphates.  The  white  lead  (q.v.)  used  in 
paintings  is  changed  by  the  hydrogen  sulphide  in  the  air  of  cities  to 
black  lead  sulphide,  Pb3(OH)2(CO3)2  +  3H2S-»3PbS  +  4H2O  + 
2C02.  This  may  be  oxidized  to  white  lead  sulphate  by  means  of 
hydrogen  peroxide: 

PbS  +  4H202  -»  PbS04  +  4H20, 

and  in  this  way  the  original  tints  of  the  picture  may  be  practically 
restored.  Organic  coloring  matters'  are  changed  into  colorless  sub- 
stances by  an  action  similar  to  that  of  hypochlorous  acid  (cf.  p.  192). 
Hence  hydrogen  peroxide  is  used  for  bleaching  silk,  feathers,  hair, 
and  ivory,  which  would  be  destroyed  by  this  more  violent  agent. 
The  products  of  its  decomposition,  being  water  and  oxygen  only, 
are  harmless,  and,  on  this  account,  it  is  used  as  a  bactericide  in 
surgery. 

Hydrogen  peroxide  exercises  the  functions  of  a  reducing  agent  in 
special  cases,  also.  Thus,  silver  oxide  is  reduced  by  it  to  silver: 

Ag20  +  H202  -*  2Ag  +  H20  +  02. 

A  solution  of  potassium  permanganate,  in  which  the  permanganic  acid 
has  been  set  free  by  an  acid:  KMnO4+H2SO4?=±  HMn04  +  KHSO4, 
is  rapidly  reduced.  The  permanganic  acid,  with  excess  of  sulphuric 
acid,  tends  to  undergo  the  first  of  the  following  changes,  provided  a 
substance,  such  as  hydrogen  peroxide,  is  present  which  can  take 
possession  of  the  oxygen  that  would  remain  as  a  balance : 

2HMnO4  4-  2H2SO4  ->  2MnS04  +  3H2O  (+  50)       (1) 

f (50)  +  5H2O2  ->  5H2O  +  502  •  (2) 

2HMnO4  +  2H2SO4  +  5H2O2  -*  2MnSO4  +  8H2O  +  5O2 

Exercises.  —  1.  What  volume  of  ozone  will  be  taken  up  by  100 
c.c.  of  water  at  12°  from  a  stream  of  oxygen  containing  7.5  per  cent 
of  ozone  (p.  103)? 

2.  At  what  temperature  will  a  ten  per  cent  solution  of  hydrogen 
peroxide  freeze  (p.  205)  ? 

3.  Write  the  thermochemical  equations  for  oxidation  of  indigo 
by  ozone  (pp.  194,  212),  and  by  hydrogen  peroxide. 


CHAPTER  XIX 
IONIZATION 

Introductory.  —  As  we  have  seen,  acids,  bases,  and  salts,  when 
dissolved  in  water,  interact  with  one  another  by  interchanging  radicals 
(p.  201).  We  have  also  learned  that  the  same  solutions  have  abnor- 
mal values  for  their  freezing-points  and  for  two  other  properties. 
These  facts  indicate  dissociation  into  the  radicals  (p.  207).  Now 
precisely  these  solutions  have  a  property  which  is  not  shared  by  any 
other  solutions,  namely,  that  of  being  conductors  of  electricity  and 
suffering  chemical  decomposition  by  the  passage  of  the  current.  Such 
solutions  are  called,  in  consequence,  electrolytes,  and  the  process  is 
named  electrolysis  (r/AeKi-pov,  and  Xvetv,  to  loosen,  i.e.  to  decompose, 
by  means  of  electricity).  Now  the  natural  inference  from  the  fore- 
going facts  is  that  the  electricity  is  carried  by  the  liberated  radicals. 
Our  first  aim  in  the  present  chapter  is  to  show  by  a  study  of  the 
chemical  changes  taking  place  in  electrolysis  that  this  inference  is  correct. 
We  then  proceed  to  discuss  the  hypothesis  of  ions,  by  means  of  which 
these  facts  are  harmonized  with  the  molecular  hypothesis.  Next, 
we  apply  the  hypothesis  to  the  explanation  of  electrolysis,  to  the 
equilibrium  between  the  ions  and  the  remaining,  undissociated  molecules, 
and  to  conductivity  phenomena  as  a  means  of  measuring  the  fraction 
ionized.  Finally  we  deduce  the  relation  between  extent  of  ionization 
and  chemical  activity. 

Incidentally,  the  facts  to  be  given  provide  the  means  of  under- 
standing the  electrolytic  processes,  many  of  them  of  great  importance 
in  chemical  industries,  to  which  frequent  reference  is  made  in  later 
chapters. 

Non- Electrolytes. — To  clear  the  ground,  we  should  first  note  the 
fact  that  only  solutions  (as  a  rule)  possess  both  of  the  properties 
in  question,  namely  that  of  conducting  and  that  of  being  decom- 
posed by  the  current.  Some  substances,  notably  the  metals  and 
materials  like  carbon,  are  conductors.  But  they  are  not  changed 

214 


ELECTROLYSIS  215 

chemically  by  the  current.  Again,  single  substances,  even  when 
they  are  such  as,  if  mixed,  yield  electrolytes,  are  not  conductors  at 
ordinary  temperatures.  Thus  hydrogen  chloride,  whether  gaseous 
or  liquefied,  is  a  nonconductor,  and  water  is  a  very  feeble  conductor, 
although  the  solution  of  the  two  conducts  exceedingly  well.  Dry 
acids,  bases,  and  salts,,  except  when  at  a  high  temperature  and  fused, 
are  likewise  nonconductors.  Furthermore,  even  amongst  solutions, 
not  all  are  conductors.  Solutions  of  sugar  and  other  substances  of 
the  same  class  (p.  205),  which  have  normal  freezing-points,  are  non- 
conductors. Only  solutions  of  acids,  bases,  and  salts  in  certain 
specified  solvents,  of  which  the  commonest  is  water,  are  electrolytes 
at  ordinary  temperatures. 

Chemical  Changes  Taking  Place  in  Electrolysis:  at  the 
Electrodes.  —  When  the  wires  from  a  battery  are  attached  to  plati- 
num plates  immersed  in  any  electrolyte  (e.g.  Fig.  21,  p.  64),  we 
observe  that  the  products  appearing  at  the  two  electrodes  are  always 
different.  They  may  be  of  several  kinds  physically,  and  will  be 
secured  for  examination  variously  according  to  their  nature.  Thus, 
when  they  are  gases  which  are  not  too  soluble,  they  may  be  collected 
in  inverted  tubes  filled  with  Tihe  solution.  Solids,  if  insoluble  in  the 
liquid,  will  either  remain  attached  to  the  electrode  or  fall  to  the 
bottom  of  the  vessel  as  precipitates.  Soluble  substances  on  the 
other  hand  will  usually  not  be  visible.  They  may  be  handled  by 
interposing  a  porous  partition  of  some  description  which  will  restrain 
the  diffusion  of  the  dissolved  body  away  from  the  neighborhood  of 
the  electrode,  while  not  interfering  appreciably  with  the  passage 
of  the  current.  Surrounding  one  electrode  with  a  porous  battery 
jar  is  a  convenient  method  for  effecting  this. 

Of  the  various  illustrations  which  we  have  encountered,  the  elec- 
trolysis of  hydrochloric  acid  (p.  64)  happens  to  have  been  the  only 
one  which  delivered  both  components  of  the  solute  with  a  minimum 
of  modification  at  the  electrodes: 

Neg.  wire,  H2  < H.C1 >  C12,  Pos.  wire. 

Hydrogen  does  not  interact  with  water,  and  chlorine  interacts  very 
slightly,  so  that  the  molecular  substances  H2  and  C12  are  promptly 
formed  from  the  elements  H  and  Cl  which  are  liberated.  The  chlo- 
rides, bromides,  and  iodides  of  those  metals  which  do  not  interact 
with  water  give  equally  simple  results: 

Neg.  wire,  Cu  < Cu.Br2 >  Br2,  Pos.  wire. 


216  COLLEGE   CHEMISTRY 

Thus,  in  electrolysis,  the  solute  seems  to  split  into  its  radicals,  and  the 
radical,  if  it  does  not  interact  with  water,  is  set  free.  A  substance 
thus  set  free  is  called  a  primary  product  of  the  electrolysis.  In  the 
foregoing  instances  both  products  are  primary. 

Usually  the  chemical  change  is  more  complex.  Thus,  when  dilute 
sulphuric  acid  is  electrolyzed,  hydrogen  and  oxygen  are  liberated  at 
the  negative  and  positive  electrodes,  respectively.  But  these  prod- 
ucts do  not  account  for  the  whole  of  the  constituents  (H2SO4).  We 
therefore  proceed  to  examine  the  materials  in  solution  round  the 
electrodes.  It  is  found  that,  as  the  action  progresses,  sul- 
phuric acid  accumulates  round  the  positive  wire,  while  the  liquid  in 
the  neighborhood  of  the  other  pole  is  gradually  depleted  of  this  sub- 
stance. In  view  of  this  fact  we  easily  explain  the  phenomenon. 
Evidently  the  substance  divides  into  its  radicals,  H  and  S04,  but 
SO4,  not  being  a  known  substance,  must  interact  with  the  water  to 
produce  sulphuric  acid  and  oxygen:  2S04  +  2H30  — »  2H2S04  +  02. 
The  whole  change  may  therefore  be  tabulated  as  follows: 

Neg.  Wire,  H2  < H2.S04 >  02  and  H2S04,  Pos.  Wire. 

Hence  the  hydrogen  is  a  primary  product,  but  the  oxygen  and  sul- 
phuric acid  are  secondary  products.  All  acids  give  hydrogen  alone  at 
the  negative  electrode,  whatever  may  be  the  product  at  the  positive. 
If  we  electrolyze  cupric  nitrate  solution,  we  obtain  a  red  deposit  of 
metallic  copper  on  the  negative  plate  and  at  the  positive  end  oxygen 
and  nitric  acid  are  formed.  We  infer,  therefore,  that  the  division 
of  the  original  molecule  was  into  Cu  and  N03,  but  that  the  latter 
interacted  with  the  water:  4N03  +  2H20 » 4HN03  -f  O2: 

Neg.  Wire,  Cu.  < Cu.(N03)2 >  02  and  HN03,  Pos.  Wire. 

With  a  solution  of  potassium  nitrate  we  find  hydrogen  and  oxygen 
appearing  at  the  negative  and  positive  electrodes  respectively.  Lit- 
mus paper,  however,  shows  the  presence  in  the  solution  of  a  base 
(potassium  hydroxide,  KOH)  at  the  negative  and  an  acid  (nitric 
acid)  at  the  positive  end.  Secondary  chemical  changes  have 
occurred  at  both  poles.  We  infer  that  the  parts  of  the  parent  mole- 
cules are  K  and  N03.  The  former,  since  it  resembles  sodium  (p.  66), 
instead  of  being  liberated,  gave  rise  to  free  hydrogen  and  potassium 
hydroxide  : 

Neg.  Wire,  H2  and  KOH  < K.NO3 >  O2  and  HN03,  Pos.  Wire. 


ELECTROLYSIS  217 

We  are  confirmed  in  these  conclusions  when  we  employ  a  pool  of 
mercury  in  place  of  the  negative  wire.  A  portion  of  the  potassium 
is  then  found  to  have  dissolved  in  the  mercury  and  escaped  inter- 
action with  the  water. 

Having  now  before  us  the  results  of  electrolyzing  some  typical 
substances,  we  bring  these  results  into  relation  with  the  facts 
described  in  the  last  chapter.  Acids  contain  hydrogen  which  pos- 
sesses certain  specific  properties  (p.  201),  and  by  electrolysis  they  all 
divide  so  as  to  give  up  this  constituent  alone  at  one  electrode.  The 
evidence  that  the  other  radical  has  different  electrical  properties 
which  carry  it  to  the  opposite  plate  is  conclusive.  Again,  salts 
undergo  double  decomposition  in  which  they  exchange  radicals  with 
acids,  bases,  and  other  salts  (p.  201),  and  we  find  that  it  is  these  very 
radicals  which  are  withdrawn  from  the  solution  by  the  influence  of  the 
electricity.  Furthermore,  the  radicals  exist  free  in  the  solution,  being 
formed  by  dissociation  of  the  molecules  (p.  207).  Hence  the  function 
of  the  electricity  seems  simply  to  consist  in  sifting  apart  the  two  kinds  of 
free  radicals  which  each  solution  contains.  It  only  remains  for  us 
to  construct  an  hypothesis  (see  below)  to  account  for  the  sifting 
action  of  the  current.  Before  turning  to  this  explanation  of  the 
phenomena,  however,  there  is  one  question  which  may  be  answered 
in  passing.  Since  a  solution  may  eventually  be  cleared  of  all  the 
hydrochloric  acid,  for  example,  which  it  contains,  we  should  like  to 
know  how  the  free  radicals  in  the  center  of  the  cell  reach  the  electrodes. 

Ionic  Migration. —  To  know  how  the  free  radicals  reach  the  elec- 
trodes, all  that  is  necessary  is  to  take  a  material,  one  (or  both)  of 
whose  radicals  is  a  colored  substance,  and  watch  the  movement  of 
the  colored  material  as  it  drifts  towards  the  electrode.  Most  salts 
which  give  colored  solutions  are  suitable.  In  very  dilute  cupric 
sulphate  solution,  for  example,  a  freezing-point  determination  shows 
that  thedepressionhas  practically  double  the  normal  value.  In  other 
words,  the  dissociation  into  the  radicals,  CuSO4«=±(Cu)  +  (S04), 
is  almost  complete.  Now,  the  blue  color  of  the  solution  cannot  be 
due  to  the  few  remaining  molecules  of  CuSO4,  for  anhydrous  cupric 
sulphate  is  colorless.  Nor  is  it  due  to  the  color  of  the  (SO4)  radicals, 
for  dilute  potassium  sulphate  and  dilute  sulphuric  acid  are  both 
colorless.  On  the  other  hand,  all  cupric  salts,  in  dilute  solution, 
have  the  same  tint.  The  color  is  therefore  that  of  the  free  cupric 


218 


COLLEGE    CHEMISTRY 


radical  (Cu).  In  order  most  clearly  to  see  the  motion  of  the  cupric 
radical,  we  place  the  cupric  sulphate  solution  in  the  middle  of  the 
space  beween  the  electrodes,  and  place  between  it  and  the  latter  a 
colorless  conducting  solution.  The  motion  of  the  blue  material 
across  the  boundary  may  then  be  easily  observed. 

The  most  convenient  arrangement  is  to  dissolve  the  cupric  sul- 
phate in  warm  water  containing  about  5  per  cent  of  agar-agar,  and 
to  fill  with  this  mixture  the  lower  part  of  a  U-tube  (Fig.  51).  The 

setting  of  the  jelly  prevents  subse- 
quent mixing  of  the  cupric  sulphate 
system  of  materials  with  the  rest 
of  the  filling  of  the  tube,  and  the 
consequent  disappearance  of  the 
boundary.  A  few  grains  of  char- 
coal may  be  scattered  on  the  sur- 
face of  the  jelly  to  mark  the  present 
limits  of  the  colored  substance, 
and  a  solution  of  some  colorless 
electrolyte,  such  as  potassium  ni- 
trate, is  added  on  each  side.  To 
prevent  agitation  of  the  liquid  by 
the  effervescence  at  the  electrodes, 
it  is  well  to  use  agar-agar  with  the 
lower  part  of  the  colorless  liquid 
also.  The  whole  is  finally  placed  in  ice  and  water,  to  prevent 
melting  of  the  jelly  by  the  heat  caused  by  resistance,  and  the  current 
is  then  turned  on. 

After  a  time,  we  observe  that  the  blue  cupric  radicals  ascend  above 
the  mark  on  the  negative  and  descend  away  from  it  on  the  positive 
side.  In  each  case  there  is  no  shading  off  in  the  tint.  The  motion 
of  the  whole  aggregate  of  colored  radicals  occurs  in  such  a  way  that, 
if  the  contents  of  the  tube  were  not  held  in  place  by  the  jelly,  we 
should  believe  that  a  gradual  motion  of  the  entire  blue  solution  was 
being  observed.  With  a  current  of  110  volts,  and  a  16-candle  power 
lamp  in  series  with  the  cell,  the  effect  becomes  apparent  in  a  few 
minutes. 

Although  the  (S04)  radicals  are  invisible,  we  may  safely  infer  that 
they  are  drifting  towards  the  positive  electrode.  Indeed,  this  can 
be  demonstrated  by  interposing  a  shallow  layer  of  jelly  containing 


FIG.  51. 


IONIZATION.  219 

some  barium  salt  a  little  distance  above  the  charcoal  layer  on  the 
positive  side.  When  the  (SO4)  reaches  this,  barium  sulphate  begins 
to  be  precipitated  and  the  layer  becomes  cloudy.  In  similar  ways 
the  progress  of  other  colorless  ions  may  be  rendered  visible. 

It  appears,  therefore,  that  electrolysis  is  not  a  local  phenomenon, 
going  on  round  the  electrodes  only,  but  that  the  whole  of  the  products 
of  the  dissociation  of  the  solute  are  set  in  motion.  It  is  on  account  of 
this  remarkable  property  of  traveling  or  migrating  towards  one  or 
other  of  the  electrodes  that  the  individual  atoms  (like  Cu),  or  groups 
of  atoms  (like  SO4),  have  been  named  ions  (Gk.  iW,  going).  The 
term  was  first  applied  by  Faraday  to  the  materials  liberated  round 
the  electrodes. 

Different  ionic  substances  move  with  different  speeds  when  pro- 
pelled by  the  same  current.  The  hydrogen  radical  of  acids  (H)  is 
the  most  speedy,  the  hydroxyl  radical  of  bases  (OH)  comes  next. 
The  actual  speeds  of  several  ions,  in  dilute  solutions  at  18°,  when 
driven  by  a  potential  difference  of  1  volt  between  plates  1  cm.  apart, 
expressed  in  cm.  per  hour  is:  H  10.8,  OH  5.6,  Cu  1.6,  S04  1.6, 
K  2.05,  Cl  2.12. 

The  Hypothesis  of  Ions.  —  That  the  molecules  of  certain  classes 
of  substances,  although  seemingly  without  chemical  interaction 
with  the  water  in  which  they  are  dissolved,  should  nevertheless  be 
decomposed  by  the  influence  of  the  water,  is  strange,  but  not  incon- 
ceivable. Heating  produces  a  somewhat  similar  effect  on  many  sub- 
stances. The  novel  fact,  for  which  an  explanation  is  demanded,  is 
that  the  molecules  of  the  products  of  the  dissociation  appear  to  be 
attracted  by  electrically  charged  plates,  which  have  been  lowered 
into  the  solution,  while  molecules  of  dissolved  sugar,  for  example, 
are  not  so  attracted.  Now  the  only  bodies  which  we  find  to  be 
conspicuously  attracted  by  electrically  charged  objects  are  bodies 
which  are  already  provided  with  electric  charges  of  their  own.  Thus 
we  are  led  to  add  to  the  molecular  hypothesis  the  assumption  that 
substances  which  undergo  dissociation  in  solution  divide  themselves 
into  a  special  kind  of  electrically  charged  molecules. 

Since  the  solution,  as  a  whole,  has  itself  no  charge,  equal  quantities 
of  positive  and  negative  electricity  must  be  produced: 

HC1  <=±  H  +  Cl     NaCl  ^  Na  +  Cl     NaOH  <=>  Na  +  OH. 


220  COLLEGE   CHEMISTRY 

This  means  that  bivalent  radicals,  on  dissociation,  will  become  ions 
carrying  a  double  charge  and  trivalent  ions  must  carry  a  triple 
charge: 

CuCl2  <±  Cu  +  2C1  CuS04  <=>  Cu  +  SO4 

K2SO4  <=±  2K  +  S=O4  FeCl3  <=»Ve++  3C1 

In  these  equations,  the  coefficients  multiply  the  charges,  as  well  as 
the  radicals  bearing  the  charges,  and  it  will  be  seen  that  the  numbers 
of  +  and— charges  produced  by  each  dissociation  are  equal.  Hence, 
univalent  ions  all  possess  equal  quantities  of  electricity,  and  other  ions 
bear  quantities  greater  than  this  in  proportion  to  their  valence.  This  is 
an  inevitable  inference  from  the  electrical  neutrality  of  all  solutions. 
It  is  confirmed  by  actual  measurement. 

To  show  that  this  hypothesis  is  adequate,  we  next  apply  it  to  the 
explanation  of  the  phenomena  of  electrolysis.  After  that  some 
seeming  objections  will  be  discussed. 

Application    to     the    Explanation    of     Electrolysis.  —  A 

battery  is  a  machine  which  maintains  two  points,  its  poles,  or  two 
wires  connected  with  them,  at  a  constant  difference  of  potential. 
One  cell  of  a  storage  battery,  for  example,  maintains  a  potential 
difference  of  two  volts.  When  the  wires  are  joined,  directly  or 
indirectly,  the  poles  are  immediately  discharged,  but  the  cell  con- 
tinuously reproduces  the  difference  in  potential  by  generating  fresh 
electricity.  Now  the  effect  of  immersing  two  plates,  one  of  which  is 
kept  by  the  battery  at  a  definite  positive  potential  and  the  other  at 
a  definite  negative  potential,  into  a  liquid  filled  with  floating  multi- 
tudes of  minute  bodies,  already  highly  charged,  may  easily  be  foreseen. 
The  figure  (Fig.  52)  will  convey  some  idea  of  the  behavior  of  the 
parts  of  a  system  such  as  we  have  imagined.  The  electrodes  are 
marked  --  and  +.  The  negatively  charged  plate  attracts  all  the 
positively  charged  particles  in  the  vessel,  and,  although  these  parti- 
cles are  in  continuous  and  irregular  motion,  they  nevertheless  begin, 
on  the  whole,  to  drift  toward  the  plate  in  question.  On  the  other 
hand,  the  negatively  charged  particles  are  repelled  by  this  plate  and 
attracted  by  the  positive  plate,  so  that  they  drift  in  the  opposite 
direction.  Those  which  are  nearest  each  plate,  on  coming  in  contact 
with  it,  will  have  their  charges  of  electricity  neutralized  by  the  oppo- 
site charge  on  the  plate,  turning  thereby  into  the  ordinary  free  forms 


O 
O 
O 
O 


©   © 


e 


O 
C? 


Ill 


IONIZATION  221 

of  the  matter  of  which  they  are  composed.  The  continuous  removal 
of  the  electrical  charges  of  the  plates  through  contact  with  ions  of 
the  opposite  charge  furnishes  occasion  for  recharging  of  the  plate 
from  the  battery,  and  thus  gives  rise  to  a  continuous  current  in  each 
wire.  Again,  the  continuous  drifting  of  positively  and  negatively 
charged  particles  in  opposite 

directions  through  the  liquid,      cathode  +  Anode 

constitutes  what,  in  the  view          —        4— cation=A§  4. 

<»        11  i  f          / — "\  anion^NOa 

of  all  external  means  of  fo 
observation,  appears  to  be 
an  electrical  current  in  the 
liquid  also.  A  magnetized 
needle,  for  example,  which  is 
deflected  when  brought  near 
to  one  of  the  wires  of  the 
battery,  is  influenced  in  the  FIG.  52. 

same  way  by  being  brought 

over  the  liquid  between  the  electrodes.  The  illusion,  so  to  speak,  of 
an  electric  current  is  complete,  although  in  reality  it  is  a  connection  of 
electricity  that  is  taking  place.  Furthermore,  the  quantity  of  elec- 
tricity being  transported  across  any  section  of  the  whole  system  is 
the  same  as  that  across  any  other,  whether  this  section  be  taken 
through  one  of  the  wires,  through  the  electrolyte,  or  even  through 
the  battery  at  any  point.  As  fast  as  the  ions  are  thus  annihilated  as 
such,  the  undissociated  molecules  (mingled  with  the  ions,  but  not 
shown  in  the  figure)  dissociate  and  produce  fresh  ones,  as  in  all 
chemical  equilibria.  Eventually,  by  continuing  the  process  long 
enough,  if  the  substances  set  free  are  actually  deposited  and  do  not 
go  into  solution  again  in  any  form,  the  liquid  can  be  entirely  deprived 
of  the  whole  of  the  solute  which  it  contains. 

The  analogy  to  the  transportation  of  a  fluid  like  water  is  notice- 
able, although  not  complete.  Water  may  be  transported  in  three 
ways.  It  may  flow  through  a  pipe,  it  may  pass  by  pouring  freely 
from  one  container  to  another,  and  it  may  be  carried  in  vessels. 
Thus  a  stream  of  water,  essentially  continuous,  might  be  arranged, 
in  which  part  of  the  passage  took  place  through  the  pipes,  part  by 
pouring  from  the  pipes  into  buckets,  and  part  by  the  carrying  of 
those  buckets  between  the  ends  of  the  pipes.  The  quantity  of  water 
passing  a  given  point  per  minute  in  this  system  would  be  the  same 


222  COLLEGE    CHEMISTRY 

at  every  part,  although  the  actual  method  by  which  the  water  was 
transported  past  the  various  points  might  be  different.  In  such  a 
disjointed  circuit  we  suppose  the  electricity  to  move  when  carried 
from  a  battery  through  an  electrolytic  cell.  It  flows  in  the  wire, 
passes  by  discharge  between  the  pole  and  the  ion,  and  is  transported 
upon  the  ions  in  the  liquid.  The  parallel  is  imperfect,  however, 
because  we  have  used  the  conception  of  two  electric  fluids  and  because 
the  ions  are  already  charged  in  the  solution,  and  before  any  connection 
with  the  battery  is  made.  They  do  not,  so  to  speak,  transport  the 
electricity  of  the  battery,  but  their  own. 

Difficulties  Presented  by  this  Hypothesis.  —  The  question 
was  raised  (p.  207),  as  to  how  we  can  imagine  separate  atoms  of 
sodium  to  exist  in  water  without  acting  upon  it,  as  the  metal  sodium 
usually  does.  But  the  ions  of  sodium  in  sodium  chloride  solution 
are  not  metallic  sodium.  They  bear  large  charges  of  electricity. 
They  possess  an  entirely  different,  and  in  fact,  by  measurement, 
much  smaller  amount  of  chemical  energy  than  free  sodium.  And, 
as  we  have  seen,  the  properties  of  a  substance  are  determined  as 
much  by  the  energy  it  contains  as  by  the  kind  of  matter.  Metallic 
sodium  and  ionic  sodium  are,  simply,  different  substances. 

We  think  of  hydrogen  chloride  and  common  salt  as  exceedingly 
stable  substances,  and  are  averse  to  believing  that  precisely  these 
compounds  should  be  highly  dissociated  by  mere  solution  in  water. 
But  it  must  be  remembered  that  in  solution  they  undergo  chemical 
change  very  easily,  and  it  is  only  in  the  dry  form  that  they  show 
unusual  stability. 

Again,  why  do  not  the  ions  combine,  in  response  to  the  attractions 
of  their  charges?  The  answer  is  that  they  do  combine,  but  the  rate 
at  which  combination  takes  place  is  no  greater  than  that  at  which 
the  molecules  decompose,  so  that  on  the  whole  the  proportion  of 
ions  to  molecules  remains  unchanged. 

Finally,  it  might  appear  that  the  assumption  that  bodies  could 
retain  high  charges  in  the  midst  of  water  is  contrary  to  all  experience. 
It  must  be  remembered,  however,  that  the  molecular,  pure  water, 
which  separates  the  ions  from  one  another,  is  a  perfect  nonconductor. 
The  moisture  which  covers  electrical  apparatus  and  causes  leakage 
of  static  electricity  is  not  pure  water,  but  a  dilute  solution  containing 
carbonic  acid  (p.  77)  and  materials  from  the  glass  of  which  the 


IONIZATION  223 

apparatus  is  made.  It  conducts  away  the  charge  electrolytically, 
by  means  of  the  ions  it  contains,  and  not  by  itself  acting  as  a 
conductor. 

Resume  and  Nomenclature*  —  The  dissociation  of  molecules 
into  ions  is  named  ionization.  The  substances  of  the  three  classes 
which  alone  are  ionized  may  be  designated  ionogens.  An  ion  may 
be  denned  as,  a  molecule  bearing  negative  or  positive  charges  of 
electricity  in  proportion  to  its  valence,  and  formed  through  the 
dissociation  of  an  ionogen  by  a  solvent  like  water. 

Each  molecule  of  the  solute  gives  two  kinds  of  ions  with  opposite 
charges.  These  two  are  forthwith  distinct  and  independent  sub- 
stances, save  that  the  attractions  of  the  charges  prevent  separation 
by  diffusion.  They  differ  from  non-ionic  substances  of  the  same 
material  composition  when  such  are  known.  The  electrical  charge 
is  one  of  the  essential  constituents,  and  when  it  is  removed  the 

properties  alter  entirely.     Thus  we  have  two  kinds  of  hydrogen, 

+ 
gaseous  molecular  hydrogen  (H2),  and  ionic  hydrogen  (H),  with 

entirely  different  chemical  properties  (p.  208). 

The  radicals  and  their  chemical  behavior  are  real,  and  all  the 
peculiarities  of  aqueous  solutions  of  acids,  bases,  and  salts  are 
experimental  facts.  Ions,  however,  like  corpuscles,  atoms,  and 
molecules,  are  part  of  our  great  system  of  formulative  hypotheses 
and  are  added  to  it  in  order  to  maintain  its  self-consistency.  Mole- 
cules are  units  which  are  not  commonly  disintegrated  by  vaporiza- 
tion (p.  87) ;  ions,  those  which  are  not  commonly  disintegrated  in 
double  decomposition  in  solution;  atoms,  those  which  are  not  com- 
monly disintegrated  in  any  chemical  action.  The  ionic  hypothesis 
was  first  suggested  by  Svante  Arrhenius,  a  Swedish  chemist,  in  1887. 

Since  the  writing  of  the  +  and  —  charges  over  the  symbols 
occupies  much  space,  we  shall  hereafter  employ  a  dot  for  the  former 
and  a  little  dash  for  the  latter:  H*,  CY,  Cu",  SO/',  Fe",  Fe"*,  NH*4. 
In  the  ions  formed  from  one  molecule,  the  number  of  dots  and 
dashes  must  be  equal. 

Since  ionic  hydrogen,  ionic  chlorine,  etc.,  are  entirely  different  in 
physical  and  chemical  properties  from  the  corresponding  free  ele- 
ments, they  should  receive  separate  names.  When  it  is  inconvenient 
to  say  "  ionic  hydrogen/'  "  ionic  nitrate  radical  "  (NOS'),  etc.,  the 
following  will  be  used: 


224 


COLLEGE   CHEMISTRY 


Sym- 

Name. 

Anion  of 

Sym- 

Name. 

Cation  of  Salts  of 

bol. 

bol. 

S04" 

Sulphate-ion 

Sulphates 

Na* 

Sodium-ion 

Sodium 

Cl' 

Chloride-ion 

Chlorides 

Fe— 

Ferric-ion 

Ferric  iron 

HSO/ 

Hydrosulphate-ion 

Bisulphates 

NH4' 

Ammonium-ion 

Ammonium 

OH7 

Hydroxide-ion 

Hydroxides 

Fe" 

Ferrous-ion 

Ferrous  iron 

(bases) 

H* 

Hydrogen-ion 

Hydrogen  (acids) 

In  using  these  terms,  note  that  sodium-ion  (with  the  hyphen)  is  the 
name  of  the  substance,  and  not  of  the  hypothetical,  charged  atom. 
When  speaking  in  terms  of  hypothesis,  therefore,  we  may  not  say 
"  a  sodium-ion,"  any  more  than  we  should  say  "  an  ionic  sodium  " 
or  "  ionic  sodiums."  To  describe  the  charged  molecule,  we  must 
write  "  a  sodium  ion,"  "  sodium  ions,"  "  chlorate  ions,"  etc. 

Faraday  distinguished  the  two  kinds  of  material  which  proceed 
with  and  against  the  positive  current  by  name.  His  terminology  is 
still  used.  Ions  which  proceed  in  the  same  direction  as  the  positive 
current  (Fig.  52)  are  called  cations  (Gk.  Kara,  down).  Such  are  H', 
Cu",  K",  NH4".  They  are  metallic  elements,  or  groups  which  play  the 
part  of  a  metal.  The  electrode  (Gk.,  6Sos,  a  path)  upon  which  they 
are  deposited,  the  negative  electrode,  is  spoken  of  as  the  cathode 
(Gk.  17  KatfoSos,  the  way  down). 

The  particles  which  move  in  the  direction  of  the  negative  current, 
and  against  that  of  the  positive,  are  named  anions  (Gk.  dm,  up). 
The  ions  Cl',  NO3',  SO/',  MnO/  are  of  this  kind.  They  are  usually 
composed  of  non-metals,  although  sometimes,  as  in  MnO/,  the  com- 
ponents may  be  partially  metallic.  They  are  set  free  at  the  positive 
electrode,  which  is  therefore  named  the  anode  (Gk.  17  avoSos,  the 
way  up).  Chemists  speak  of  metals  and  non-metals  as  positive  and 
negative  elements,  respectively  (cf.  p.  82),  even  when  electrical 
relations  are  not  directly  in  question,  and  ions  are  not  concerned. 

Applications :  Ionic  Equilibrium.  —  Since  the  ions  are  chemi- 
cally different  from  their  parent  molecules,  their  formation  represents 
a  variety  of  chemical  change.  The  change  does  not  involve  any 
chemical  interaction  with  the  water.  It  is  simply  a  dissociation, 
i.e.  reversible  decomposition  of  the  dissolved  substance. 

From  the  fact  that  the  proportion  of  molecules  ionized  is  shown  to 


IONIZATION 


225 


become  greater  as  more  and  more  of  the  solvent  is  added  (p.  206), 
and  that  removal  of  the  solvent  diminishes  the  proportion  of  ions  to 
molecules,  and  finally  leaves  us  the  substance  entirely  restored  to 
the  molecular  condition,  we  know  that  this  is  a  reversible  action  and 
therefore  a  true  dissociation.  The  molecules  and  their  ions  adjust 
themselves  like  the  constituents  in  any  case  of  chemical  equilibrium 
(pp.  174-180) : 

NaCl  <=±  Na'  +  Cl'. 

The  chemical  behavior  of  substances  in  ionic  equilibrium  will  be 
discussed  in  the  next  chapter  (see  p.  234). 

*  The  mode  of  formulation  previously  used    (p.  180)  may  be 


FIG.  53. 

employed  here.  If  [NaCl],  [Na*],  and  [Cl']  stand  for  the  molecular 
concentrations  (numbers  of  moles  per  liter)  at  equilibrium  of  the 
molecules,  and  the  two  ions,  respectively,  we  have  an  equilibrium 
constant  (cf.  p.  181),  in  this  case  called  the  ionization  constant: 

[Na']  X  [Cl'] 
[NaCl] 


K  = 


When  we  dissolve  a  single  substance  which  gives  only  two  ions,  the 
molecular  concentrations  of  the  ions  are  necessarily  equal.  When 
some  other  ionogen  with  a  common  ion  is  present,  however,  the 
values  of  [Na']  and  [Cl']  will  be  different. 

*  This  paragraph  is  not  required  for  understanding  anything  that  immedi- 
ately follows  (see  Chap,  xxxiv). 


226  COLLEGE   CHEMISTRY 

Applications:  To  the  Interpretation  of  Conductivity 
Measurements.  —  We  have  seen  that  when  the  solution  of  an 
ionogen  is  diluted,  the  proportion  of  ions  to  undissociated  molecules 
increases,  while  removal  of  a  part  of  the  solvent  has  the  opposite 
effect  (p.  224).  Now,  a  change  in  the  number  of  ions  naturally 
modifies  the  capacity  of  the  liquid  for  carrying  electricity,  so  that 
observation  of  the  changes  in  the  conductivity  of  a  solution,  when 
the  concentration  is  altered,  supplies  the  simplest  means  of  studying 
the  phenomena  of  ionization. 

A  glass  trough  and  amperemeter  *  (Fig.  53)  may  be  used  to  illus- 
trate this  principle.  The  electrodes  are  long  strips  of  copper  foil, 
which  pass  down  at  the  ends  of  the  trough.  After  placing  the  two 
instruments  in  circuit  with  a  source  of  electricity,  we  first  pour  very 
pure  water  into  the  cell.  With  this  arrangement,  the  amperemeter 
does  not  indicate  the  passage  of  any  current  of  electricity.  Con- 
centrated (36%)  hydrochloric  acid  is  next  cautiously  added  through 
a  long-stemmed  dropping  funnel,  so  that  it  forms  a  shallow  layer 
below  the  water,  and  the  funnel  is  withdrawn.  The  situation  at  this 
stage  is  that  a  definite  amount  of  hydrogen  chloride  dissolved  in  a 
small  amount  of  water  fills  what  was  before  a  gap  in  the  electric 
circuit.  The  deflection  of  the  needle  in  the  amperemeter  indicates 
that  a  certain  current  of  electricity  is  able  to  pass  through  this  acid. 
When  we  now  stir  the  surface  of  the  acid  very  gently  with  a  thin 
glass  rod,  the  amperemeter  instantly  responds,  showing  an  increase 
in  conductivity.  As  we  stir,  the  conductivity  increases,  and  the 
increase  ceases  only  when  the  liquid  has  become  homogeneous. 
Introduction  of  an  additional  supply  of  water  will  improve  the  con- 
ductivity still  more,  but  the  effect  becomes  less  and  less,  until  no 
change  on  farther  dilution  is  perceptible.  Reasoning  about  these 
effects,  we  perceive  that  the  amount  of  hydrochloric  acid  has  not 
altered  during  the  experiment.  Yet  the  quantity  of  conducting 
material  between  the  electrodes  must  have  become  greater,  for  the 
carrying  power  of  the  whole  has  improved.  We  were  therefore 
observing  the  progress  of  a  chemical  change  of  the  nonconducting 
hydrogen  chloride  into  a  conducting  material.  In  terms  of  the 
hypothesis,  hydrogen  chloride  molecules  do  not  carry  electricity 

*  An  amperemeter  of  low  resistance,  0.5-1  ohm,  must  be  used.  The  current 
(direct,)  passing  from  the  dynamo  through  one  (or  two)  32-candle  lamps, 
after  being  reduced  to  8-10  volts  by  means  of  a  suitable  shunt,  may  be  em- 
ployed. 


IONIZATION  227 

(p.  214),  but  the  hydrogen  and  the  chloride  ions  into  which  it  was 
gradually  altered  by  chemical  change  do.  Furthermore,  the  change 
practically  ceased  at  great  dilution,  for  the  dissociation  into  ions 
was  then  practically  complete.  If  we  could  conveniently  have 
started  with  only  liquefied,  dry  hydrogen  chloride  in  the  cell,  we 
should  have  observed  the  whole  range  of  changes  from  zero  to  the 
maximum. 

When  a  saturated  solution  of  cupric  chloride  is  used  instead  of 
hydrochloric  acid,  dilution  is  accompanied  by  a  similar  improve- 
ment in  conductivity.  Here  we  notice,  besides,  that  the  yellowish- 
green  liquid,  with  which  we  start,  changes  to  a  pale  blue,  as  the 
yellowish-brown  molecules  of  cupric  chloride  are  dissociated  and  the 
color  of  the  solution  becomes  more  exclusively  that  of  the  copper 
ions.  When  the  solution  has  become  perfectly  blue,  further  dilution 
is  seen  to  affect  the  conductivity  but  slightly. 

Reasoning  still  further  about  these  phenomena  we  see  that,  if  we 
start  with  a  fixed  amount  of  a  given  substance,  the  conductivities  at 
different  stages  of  the  dilution  must  be  proportional  to  the  numbers  of 
ions,  and  the  maximum  conductivity  attainable  by  great  dilution 
must  represent  the  effect  when  the  whole  material  has  become  ionic. 
Thus,  if  the  conductivity  at  the  maximum  is  represented,  say,  by  5, 
then  at  the  dilution  where  the  conductivity  is  2,  the  proportion  of 
the  whole  which  is  ionized  is  2/5.  When  the  conductivity  becomes 
4,  4/5  of  the  molecules  are  dissociated  and  the  degree  of  ionization 
is  .8.  When  the  conductivity  becomes  5,  5/5,  or  all,  of  the  molecules 
are  dissociated.  For  example,  in  hydrochloric  acid,  if  we  take  the 
normal  solution  (p.  99)  containing  36.5  g.  of  acid  per  liter  as  the 
unit  of  concentration,  the  fractions  ionized  at  various  concentrations 
are  as  follows:  10AT,  0.17;  N,  0.78;  N/W,  0.91;  N/lOO,  0.96. 
Thus,  measurements  of  conductivity  enable  us  to  study  the  ionic 
decomposition  of  all  ionogens,  and  to  state  accurately  the  fraction 
ionized,  at  each  concentration,  in  solutions  of  every  ionogen.  This 
information  is  obviously  most  valuable,  for  it  places  us  in  a  position 
to  know  the  exact  constitution  of  every  solution  we  use  in  the 
laboratory.  In  the  following  section  the  data  on  which  such  knowl- 
edge can  be  based  is  given.  In  the  next  chapter  the  mode  of 
applying  the  data  is  explained. 

Constitution  of  Solutions  of  Ionogens :  Fractions  Ionized.— 

The  dilute  acids  used  in  the  laboratory  are  generally  of  six  times 


228 


COLLEGE   CHEMISTRY 


normal  (6A7)  concentration.  But  often  we  add  only  a  drop  or  two 
to  a  large  bulk  of  liquid,  so  that  the  acids  are  commonly  very  dilute 
as  actually  employed.  The  solutions  of  salts  are  of  different  strengths, 
but  the  great  majority  are  of  normal  (N),  or  even  smaller  concentra- 
tion. In  practice  they,  also,  are  still  further  considerably  diluted 
before  use.  If,  therefore,  we  give  the  fractions  ionized  (total  mole- 
cules =  1)  in  decinormal  solutions  (except  where  otherwise  specified), 
the  reader  will  be  able  to  estimate  roughly  the  proportion  of  each 
kind  of  ions  in  any  application  of  the  reagent.  In  the  case  of  acids 
containing  more  than  one  displaceable  hydrogen  unit,  the  kind  of 
ionization  on  which  the  figure  is  based  is  indicated  by  a  period. 
Thus  H.HCO3  means  that  the  whole  of  the  ionization  is  assumed  to 
be  into  H*  and  HC03'. 

FRACTION  IONIZED. 
ACIDS. 


Nitric  acid      

0.92 

Carbonic  acid,  H.HCO3     .    .  0.0017 

Nitric  acid  (cone.,  62%)    . 
Hydrochloric  acid    .... 

0.096 
0.91  - 

Carbonic       acid,       H.HCO3 
(AT/25)     0  0021 

Hydrochloric     acid     (cone. 
35%)     
Sulphuric  acid,  H.H.SO4   . 

0.136 
0.58 

Hydrogen  sulphide,  H.HS    .0,0007 
Boric  acid,  H.H2BO3  .    .    .    .   0.0001 
Hydrocyanic  acid    .    .               0  0001 

Sulphuric  acid  (cone.,  95%) 
Hydrofluoric  acid        .    .    . 

0.01 
0  15 

Permanganic  acid  (AT/2)    .    .   0.93 
Hydriodic  acid  (AT/2)     .    .    .   0.90 

Oxalic  acid,  H.HC2O4     .    .    . 
Tartaric  acid,  H.HT  .... 
Acetic  acid  (N)     
Acetic  acid     

0.50 
0.08 
0.004 
0.013 

Hydrobromic  acid  (AT/2)  .    .   0.90 
Perchloric  acid  (N  /2)     .    .    .   0.88 
Chloric  acid  (AT/2)   0.88 
Phosphoric  acid,  H.H2PO4    .0.26 

BASES. 


Potassium  hydroxide 
Sodium  hydroxide 
Barium  hydroxide 
Lithium  hydroxide  (AT) 
Ammonium  hydroxide 
Tetramethylammonium   hy 
droxide  (AT/16)     .... 

Potassium  chloride  .  .  . 
Potassium  nitrate  .... 
Potassium  acetate  .... 
Potassium  sulphate  .  .  . 
Potassium  carbonate  .  .  . 
Potassium  chlorate  .  .  . 
Ammonium  chloride  .  .  . 
Sodium  chloride  (AT)  .  .  . 
Sodium  chloride  (AT/2)  .  . 

Sodium  chloride 

Sodium  nitrate 

Sodium  acetate 

Sodium  sulphate 


ie     . 

.    .  0.89 
.    .   0.84 
.    .   0.80 

Strontium  hydroxide  (AT  /64)  0.93 
Barium  hydroxide  (AT  /64)  .  0.92 
Calcium  hydroxide  (AT  /64)  .0.90 

(AT) 
ide  . 

.    .   0.63 
.   0.014 

Silver  hydroxide  (AT/1783)  .0.39 
Water  .  .0.0,1 

0.96 


SALTS. 


0.86 
0.83 
0.85 
0.71 
(0.70) 
0.82 
0.85 
0.67 
0.73 
0.84 
0.83 
0.78 
0.69 


(0.52) 


Sodium  bicarbonate, 

Na.HCO3  (AT)     ... 
Sodium  phosphate, 

Na2.HPO4  (AT/32)     .    .    .    .0.83 
Sodium  tartrate  (AT /32)     .      (0.78) 

Barium  chloride 0.76 

Calcium  sulphate  (AT /100)     .  0.63 

Cupric  sulphate 0.38 

Silver  nitrate 0.81 

Zinc  sulphate 0 . 39 

Zinc  chloride 0 . 73 

Mercuric  chloride     .    .    .      (<0.01) 
Mercuric  cyanide Minute 


.  ttf  ..Jt 

IONIZATION  229 


In  addition  to  their  use  in  showing  the  nature  of  the  reagents 
employed  in  the  laboratory  (p.  227),  these  numbers  show  also  to 
what  extent  any  pair  of  ionic  substances  will  unite  when  mixed 
(see  pp.  233-234),  and  they  likewise  indicate  the  chemical  activity 
of  the  ionogens  when  in  solution  (see  next  section). 

Relation  of  lonization  to  Chemical  Activity.  —  These 
tables  may  be  used  for  reference.  The  import  of  the  following 
general  statements,  drawn  from  the  tables,  should  be  memorized: 

1.  Salts,  with  the  exception  of  those  of  mercury,  are  all  well 
ionized.     In  actions  involving  their  ions,  salts  are  therefore  all  of  the 
same  order  of  activity,  for  a  dilute  solution  of  every  salt  contains  a 
large  amount  of  the  ionic  components. 

2.  Acids  show  the  most  extreme  differences  in  their  degrees  of 
ionization.     That  is  to  say  their  solutions  must  contain  very  different 
concentrations    of    hydrogen-ion.     Since    their    activity    as    acids 
depends  on  this  substance  (p.  208),  and  the  activity  of  a  substance  is 
proportional  to  its  concentration  (p.  180),  it  follows  that  acids  will 
show  very  great  differences  in  apparent  chemical  activity.     At  this  point, 
therefore,  we  emerge  from  semi-physical  discussion  of  the  subject  and 
reach  something  of  definite,  practical  application  in  chemical  work. 

The  data  show  that  acids  may  be  divided  roughly  into  four  classes 
with  different  degrees  of  acidic  activity: 

(a)  The  ionization  in  decinormal  solution  exceeds  70  per  cent;  e.g. 
nitric  acid  and  hydrochloric  acid.  These  are  the  acids  which  are 
chemically  most  active,  for  their  solutions  contain  a  relatively  high 
concentration  of  hydrogen-ion. 

(6)  The  ionization  is  between  70  and  10  per  cent;  e.g.  sulphuric 
acid  and  phosphoric  acid.  These  acids  are  noticeably  less  active, 
for  their  solutions  contain  a  lower  concentration  of  hydrogen-ion. 

(c)  The  ionization  is  between  10  and  1  per  cent;  e.g.  acetic  acid. 
These  are  the  weaker  acids,  for  their  solutions  contain  a  very  small 
concentration  of  hydrogen-ion. 

(d)  The  ionization  is  less  than  1  per  cent;  e.g.  carbonic  and  boric 
acids.    These  are  the  feeble  acids,  for  their  solutions  contain  only 
a  minute  concentration  of  hydrogen-ion. 

3.  The  bases  show  two  classes: 

(a)  lonization  high;  e.g.  potassium  hydroxide.  These  bases  are  ac- 
tive, for  their  solutions  contain  a  high  concentration  of  hydroxide-ion. 


230  COLLEGE    CHEMISTRY 

(b)  lonization  less  than  2  per  cent;  e.g.  ammonium  hydroxide. 
These  bases  are  weak  on  account  of  the  low  concentration  of 
hydroxide-ion. 

4.  Water  is  less  ionized  than  any  other  substance  in  the  list.  It 
shows  therefore,  as  we  already  know,  usually  little  or  no  interaction 
with  acids,  bases,  or  salts,  and  hence  is  valuable  as  a  solvent  for 
these  substances.  Its  ions  are  H'  and  OH',  and  it  is  thus  as  much 
an  acid  as  a  base. 

Exercises.  —  1.  With  solutions  of  the  following  substances, 
state,  (a)  what  will  be  the  products  of  electrolysis,  (6)  whether  each 
is  primary  or  secondary,  and  (c)  how  they  may  be  isolated  in  each 
case :  Potassium  chlorate,  potassium  iodide,  potassium  iodate,  silver 
sulphate,  sodium  peroxide. 

2.  Make  equations   (p.  220)  showing  the  ionic  and  molecular 
materials  in  solutions  of  potassium  bromide,  potassium  bromate, 
sodium  periodate,   aluminium  chloride,  zinc  sulphate.     Mark  the 
charges  on  the  ions  and  give  the  name  of  each  ionic  substance  (p.  224). 

3.  Prepare  lists  of  other  anions  and  cations  which  have  been 
encountered,  giving  the  formula  and  number  of  charges  of  electricity 
in  each  case. 

4.  If  the  conductivity  of  sodium  chloride  solution  at  the  maximum 
is  110,  and  at  greater  concentrations  is  as  follows:  N,74A]  AT/10,  92.5; 
AT/100,  103,  calculate  the  fraction  ionized  at  each  concentration. 

5.  If  the  conductivity  of  acetic  acid  solution  at  the  maximum  is 
352,  and  at  greater  concentrations  is  as  follows:  ION,  0.05;  N,  1.32; 
AT/10,  4.6;  AT/100,   14.3,  calculate  the   fraction   ionized   at   each 
concentration. 

6.  If  1  c.c.  of  dilute  hydrochloric  acid  (GAT)  is  added  to  30  c.c. 
of  an  aqueous  solution,  what  is  the  reacting  concentration  of  the 
acid? 

7.  Classify  all  the  acids  in  the  table  (p.  228)  according  to  the  four 
classes  (p.  229). 


CHAPTER  XX 
IONIC    SUBSTANCES    AND    THEIR   INTERACTIONS 

IN  this  chapter,  after  enumerating  the  various  classes  of  ionogens, 
and  the  various  kinds  of  ionic  substances,  we  discuss  the  interactions 
of  the  latter.  We  consider  first  the  relations  of  the  ionic  and  the 
molecular  substances  (in  equilibrium) when  a  single  ionogen  is  present, 
and  then  take  up  the  ways  in  which  such  an  ionic  equilibrium  is  dis- 
placed. Finally  we  discuss  some  of  the  useful  ionic  interactions,  in 
which  the  equilibria  are  displaced  so  far  that  practically  complete 
interaction  occurs:  namely,  precipitation,  neutralization,  and  dis- 
placement. 

The  Classes  of  Ionogens :  Mixed  Ionogens  and  Double 
Salts.  —  Acids  are  classified  according  to  the  number  of  hydrogen 
units  in  their  molecules.  Thus  chloric  acid  HC103  is  a  monobasic 
acid,  sulphuric  acid  H2S04  a  dibasic  acid,  and  phosphoric  acid  H3PO4 
a  tribasic  acid.  These  terms  relate  to  the  fact  that,  in  neutralization 
(pp.  189,  191)  the  acids  interact  with  one,  two,  or  three  molecules 
of  a  base  like  sodium  hydroxide. 

Bases  are  named  in  a  similar  way:  sodium  hydroxide  NaOH  is  a 
monoacid  base,  calcium  hydroxide  Ca(OH)2  is  a  diacid  base. 

Salts  like  KC1  and  Na2C03  are  neutral  (see  acid  salts,  below)  or 
normal  salts,  and  NaKCO3  and  Ca(OCl)Cl  (bleaching  powder,  p.  189) 
are  mixed  salts. 

The  most  interesting  classes  of  mixed  salts  are  the  acid  salts 
(p.  171)  and  the  basic  salts.  In  acid  salts,  like  NaHSO4  (p.  117)  and 
KH2PO4  (p.  118),  all  the  hydrogen  of  the  acid  has  not  been  replaced 
by  a  metal.  In  basic  salts,  like  Ca (OH)C1,  part  of  the  basic  hydroxyl 
remains. 

There  are  also  many  double  salts,  like  ferrous-ammonium  sulphate 
(NH4)2SO4,FeSO4,6H2O,  and  alum  (q.v.),  some  of  which  are  in 
common  use. 

All  these  substances  are  ionogens  (p.  223).  The  mixed  and  double 
salts  are,  naturally,  dissociated  into  more  than  two  ionic  substances, 

231 


232  COLLEGE   CHEMISTRY 

The  Kinds  of  Ionic  Substances  Furnished  by  lonogens.  — 

The  mode  of  naming  ionic  substances  has  already  been  given  (p.  224). 

Acids,  when  dissolved  in  water,  all  furnish  hydrogen-ion  H*  and 
a  negative  ionic  substance  (anion),  e.g.,  H.CL,  H2.SO4.  The  period 
separates  the  ions,  and  shows  therefore  their  composition  in  each 
case.  The  solutions  differ  from  those  of  salts  in  the  constant  pres- 
ence of  hydrogen-ion,  and  in  the  absence  of  any  other  positive  ion. 

Hydrogen-ion  is  a  colorless  substance.  It  is  sour  in  taste,  and  its 
presence  is  recognized  by  the  fact  that  it  turns  blue  litmus  red  (see 
Indicators,  below).  These  properties  serve  as  tests  for  acids,  as 
they  are  not  interfered  with  by  other  ionic  substances  which  may  be 
present.  Hydrogen-ion  is  univalent  and,  when  combined  with 
negative  radicals  of  salts,  gives  the  (molecular)  acids.  The  activity 
of  acids  depends  upon  the  concentration  of  the  hydrogen-ion  they 
furnish  (pp.  208,  229),  and  therefore  upon  their  solubility  and  the 
degree  of  ionization  of  the  dissolved  molecules.  Some  furnish  so 
little  hydrogen-ion  that  their  action  on  litmus  can  hardly  be  detected. 

Bases  all  furnish  hydroxide-ion  OH'  and  some  positive  ionic  sub- 
stance (cation),  K.OH,  NH4.OH,  Zn.(OH)2.  Their  solutions  differ 
from  those  of  salts  in  the  constant  presence  of  hydroxide-ion  and  in 
the  absence  of  any  other  anion.  The  more  active  bases,  that  is 
those  which  are  soluble  and  highly  dissociated,  so  that  they  give  a 
high  concentration  of  hydroxide-ion,  are  called  alkalies.  Such  are 
potassium  and  sodium  hydroxides.  They  are  often  named  caustic 
alkalies  and,  individually,  caustic  potash  and  caustic  soda.  The 
solutions  are  called  lyes. 

Hydroxide.-ion  is  a  colorless  substance.  Properties  which  serve  as 
tests  for  bases  are  that  hydroxide-ion  possesses  a  soapy  taste  and 
turns  red  litmus  blue  (see  Indicators,  below).  It  is  univalent,  and 
combines  with  positive  radicals  to  form  (molecular)  bases. 

Salts  furnish  positive  and  negative  ionic  substances,  which  may  be 
either  simple  or  composite,  Na.Cl,  Na.NO3,  NH4.C1,  NH4.NO3. 
Some  ionic  substances  are  colored,  Cu**  (cupric-ion)  blue,  Cr"*  red- 
dish-violet, Co"  pink,  MnO/  (permanganate-ion)  purple,  O2O7" 
(dichromate-ion)  orange,  but  most  of  them  are  colorless,  K",  Na', 
Zn",  Cl',  I',  NO3',  SO/.  They  vary  in  taste,  some  being  salt,  some 
astringent,  some  bitter.  The  ionic  materials  characteristic  of  salts 
do  not  affect  litmus,  and  individual  tests  are  required  for  each. 
Usually  we  add  some  other  ionic  substance,  with  which  the  ion 


IONIC   SUBSTANCES  233 

thought  to  be  present  combines  to  form  an  insoluble,  molecular 
substance  of  known  color,  or  appearance,  and  examine  the  precipitate 
if  any  appears.  Thus,  when  the  presence  of  chloride-ion  CY  is  sus- 
pected, we  may  add  a  solution  containing  silver-ion  Ag',  expecting 
to  obtain  a  precipitate  of  silver  chloride  AgCl  (CY  +  Ag*  — >  AgCl|). 
In  dilute  solutions  of  salts,  the  ions  are  almost  always  numerous  in 
comparison  with  the  molecules  (p.  229) ,  so  that  salts  are  practically 
all  active  and  their  solutions  almost  always  respond  readily  to  the 
tests  for  the  ions  they  contain.  The  art  of  detecting  the  various 
ionic  substances  present  in  a  solution  constitutes  a  large  part  of  the 
branch  of  chemistry  called  qualitative  analysis. 

All  the  known  ionic  substances  are  found  in  solutions  of  salts. 
The  only  ions  which  are  not  characteristic  of  salts,  although  some- 
times occurring  in  their  solutions  (see  acid  and  basic  salts,  above), 
are  hydrogen-ion  H*,  and  hydroxide-ion  OH'. 

It  will  assist  the  reader  if  the  following  facts  are  kept  in  mind. 
The  elements  which  can  form  a  simple  positive  ion  are  the  metallic 
elements  (p.  82,  and  see  Chaps,  xxiii  and  xxxii).  Non-metallic 
elements,  like  nitrogen,  may  be  present  in  a  positive  ion,  as  in  NH4", 
but  never  exclusively.  In  other  words,  we  know  no  such  substances 
as  nitrogen  sulphate,  or  carbon  nitrate.  Conversely,  the  metals 
are  frequently  found  in  the  negative  ion,  but  never  constitute  it 
exclusively.  They  are  then  usually  associated  with  oxygen,  as  in 
MnO/,  and  Cr2O7". 

The   Ionic    Equilibrium  ivitli  a  Single  lonogen.  —  In  the 

ionization  of  a  molecular  substance,  the  chemical  change  is  incom- 
plete and  the  system  reaches  a  condition  of  equilibrium  (p.  225). 
The  action  is,  therefore,  reversible,  and  there  are  thus  two  routes  to 
the  same  equilibrium  point.  This  fact  must  not  be  forgotten,  for 
we  have  to  consider  the  union  of  ionic  substances  even  more  often 
than  the  converse  change.  Now,  the  degrees  of  ionization  of  various 
ionogens  tell  us  the  location  of  the  equilibrium  point,  and  therefore 
the  extent  of  the  chemical  change  involved  in  reaching  this  point  by 
either  route,  that  is,  either  by  the  dissociation  of  molecules  or  by  the 
union  of  ions.  In  a  class  of  interactions,  of  which  all  are  incomplete, 
and  only  those  are  interesting  and  useful  which  approach  complete- 
ness, we  require  some  means  of  knowing  which  are  complete  and 
why  they  are  so.  The  table  of  fractions  ionized  (p.  228)  supplies 
most  of  the  required  information. 


234  COLLEGE   CHEMISTRY 

i 

To  illustrate,  take  the  case  of  a  single  ionogen.  When  we  place 
hydrogen  chloride  in  decinormal  solution,  0.91  of  the  molecules  dis- 
sociate. Conversely,  when  we  start  with  the  hydrogen-ion  and 
chloride-ion,  say  by  mixing  two  solutions  each  of  which  contains  one 
of  them,  then  1-0.91,  or  only  0.09  of  these  ionic  substances  will 
combine. 

This  exemplifies  the  case  of  an  active  acid.  The  following  equa- 
tions show  the  data  for  six  typical  substances,  namely,  two  acids, 
two  bases,  and  two  salts: 

(  9%)HC1   «=±H'  +  Cl'(91%),      (98.7%)HC2H302<r»IT    +  C2H3O2'(1. 
(11%)KOH«=>  K'  +OH'(89%),     (98.6  %)NH4OH   <=±NH4'    +    OH' 

(16%)NaCl«=±Na'  +  Cl/(84%),         (62%)CuSO4      <=±Cu"    +  SO/'  (38 


These  samples  are  chosen  to  illustrate,  in  each  pair,  the  extremes. 
Thus,  with  potassium-ion  and  hydroxide-ion  little  union  takes  place, 
while  with  ammonium-ion  and  hydroxide-ion  the  union  is  practically 
complete.  In  the  case  of  the  soluble  salts,  however,  there  are  almost 
(p.  228)  no  cases  of  considerable  union  of  the  ions  in  dilute  solutions. 
The  case  of  water  is  one  of  the  most  extreme  : 

(99.95%)  H20<=±H*  +  OH'  (0.041%). 

Hydroxide-ion  and  hydrogen-ion  thus  unite  almost  completely. 

Similar  reasoning  enables  us  to  handle  the  more  complex,  but  very 
common  case  of  the  mixing  of  two  ionogens.  The  degrees  of  ioniza- 
tion  tell  us  the  exact  condition  of  each  system  separately,  before 
mixing.  The  result  of  the  mixing  is  best  understood  by  viewing  the 
change  as  consisting  in  a  displacement  of  each  of  the  equilibria  by 
the  action  of  the  components  of  the  other.  We  consider,  therefore, 
next,  the  displacement  of  ionic  equilibria. 

The  Displacement  of  Ionic  Equilibria.  —  Equilibria  are 
displaced  by  changes  which  favor  or  disfavor  one  of  the  opposed 
actions  (p.  177).  There  may  be  either,  (1)  a  physical  change  in  the 
conditions,  or  a  chemical  interaction  which  (2)  adds  to,  or  (3) 
removes  one  of  the  interacting  substances.  Each  of  these  may  be 
illustrated  in  turn. 

As  an  example  of  the  first,  we  have  the  effect  of  changing  the  amount 
of  the  solvent  (p.  226).  Adding  more  of  the  solvent  reduces  the  con- 
centration of  the  ionic  materials  and  disfavors  their  union,  so  that 
it  indirectly  promotes  dissociation.  On  the  other  hand,  evaporat- 


IONIC   SUBSTANCES  235 

ing  off  a  part  of  the  solvent  favors  the  encounters  of  the  ions  and 
promotes  combination.  When  the  solvent  is  at  last  entirely  gone, 
the  whole  material  is  molecular  (p.  225). 

In  cases  where  the  ionic  and  molecular  substances  are  all  colorless, 
these  changes  can  be  followed  only  by  a  study  of  the  freezing-points 
or  other  similar  properties  of  the  solutions  (p.  206).  But  when  the 
substances  are  of  different  colors,  the  changes  can  also  be  seen. 
Thus,  cupric  bromide  in  the  solid  form  is  a  jet  black,  shining,  crystal- 
line substance.  When  treated  with  a  small  amount  of  water  it 
forms  a  solution  which  is  of  a  deep  reddish-brown  tint,  giving  no 
hint  of  resemblance  to  a  solution  of  any  cupric  salt.  This  doubtless 
represents  the  color  of  the  molecules.  When  more  water  is  added, 
the  deep  brown  gives  place  gradually  to  green,  and  finally  to  blue. 
The  latter  is  the  color  of  the  cupric-ion  (Cu"),  and  is  familiar  in  all 
solutions  of  cupric  salts.  The  colorless  nature  of  solutions  of  potas- 
sium and  sodium  bromides  shows  that  bromide-ion  (Br')  is  without 
color.  Hence,  in  the  present  instance  it  is  invisible.  We  are  thus 
watching  the  forward  displacement  of  the  equilibrium: 

CuBr2  (brown)  ±^  Cu"  (blue)  +  2Br'. 

If  1  g.  of  the  solid  is  taken,  it  dissolves  in  about  its  own  weight  of 
water,  and  independent  measurement  shows  that  there  is  relatively 
little  ionization.  Hence  the  solution  is  deep  brown.  When  10  c.c. 
of  water  has  been  added,  70  per  cent  of  the  salt  is  ionized,  and  the 
solution  is  green.  With  40  c.c.  of  water,  only  19  per  cent  remains  in 
molecular  form,  and  the  blue  color  of  the  cupric-ion  entirely  overbears 
the  tint  of  the  molecules.  If  we  now  remove  the  water  by  evapora- 
tion, all  these  changes  are  reversed.  When  30  c.c.  of  the  water  has 
been  driven  off,  the  solution  is  green.  As  the  evaporation  of  the 
remaining  10  c.c.  progresses,  the  brown  color  appears.  When  the 
water  is  all  gone,  the  black  residue  remains.  Here  we  are  observing 
the  backward  displacement  of  the  equilibrium,  CuBr2  ±5  Cu"  4-  2Br'. 
2.  Cupric  bromide  may  be  used  to  illustrate  also  the  chemical 
methods  of  displacing  equilibria.  Thus,  we  may  show  the  effect  of 
adding  more  of  one  of  the  reacting  substances.  If,  at  the  green  stage, 
we  dissolve  solid  potassium  bromide  in  the  liquid  (KBr  <=»  K*  +  Br'), 
the  increased  concentration  of  bromide-ion  causes  more  vigorous 
interaction  of  the  ions,'  and  the  molecules,  with  their  brown  color, 
become  prominent  again.  Adding  cupric  chloride  increases  the  con- 


236  COLLEGE   CHEMISTRY 

centration  of  cupric-ion  and  has  the  same  effect.  In  either  case, 
renewed  dilution  with  water  reduces  the  concentrations  of  all  the 
ions  once  more,  the  molecules  become  fewer,  and  the  brown  color  is 
displaced  by  the  blue  for  the  second  time. 

3.  Finally,  the  displacement  of  the  same  equilibrium  by  removing 
one  of  the  interacting  substances  may  be  illustrated.  Thus,  if  the 
chocolate-brown  solution,  in  which  molecular  cupric  bromide  pre- 
dominates, is  shaken  with  pulverized  lead  nitrate  (and  filtered),  two 
changes  are  noticed.  A  pale  yellow  precipitate  of  lead  bromide 
appears  (Pb"  +  2Br'  — >  PbBr,),  and  the  brown  color  fades  into 
green.  Here  the  displacement  is  the  opposite  of  the  last.  Instead 
of  reinforcing  one  of  the  ions,  we  have  reduced  the  concentration, 
and  in  fact  almost  entirely  removed  one  of  them.  This  has,  naturally, 
stopped  the  interaction  of  the  Cu"  and  Br'  which  reproduces  the 
brown,  molecular  CuBr2.  Hence  the  dissociation  of  the  latter  has 
continued  to  exhaustion  of  the  whole  molecular  material. 

The  reader  will  find  that  the  behavior  of  these  ionic  equilibria,  and 
the  way  in  which  we  discuss  and  explain  it,  are  complete  parallels 
of  the  behavior  and  explanation  in  the  case  of  ordinary  equilibria 
(pp.  91,  176),  which  should  now  be  reexamined.  The  illustrations 
in  the  present  section,  and  particularly  the  third  (cf.  p.  183),  should 
be  considered  until  every  feature  is  perfectly  clear.  They  furnish  the 
key  to  understanding  the  applications  which  follow.  One  fact  must 
not  escape  notice,  and  that  is  that  in  none  of  the  three  instances  was 
the  forward  action  (the  dissociation)  in  itself  affected.  The  molecules 
of  cupric  bromide  have,  as  we  should  expect,  a  certain  tendency  to 
decompose.  No  encounters  between  these  molecules  are  required 
for  mere  decomposition.  Hence  their  decomposition  is  not  influenced 
by  their  nearness  to,  or  remoteness  from  one  another  (illustration  1), 
nor  by  the  presence  of  any  other  kinds  of  molecules  or  ions  (illustra- 
tions 2  and  3).  The  effect,  whether  it  involved  an  apparent  increase, 
or  a  diminution  of  the  dissociation,  was  always  accomplished  by 
altering  the  concentration  of  the  ionic  substances,  and  therefore  the 
activity  of  the  reverse  action. 

Applications  :  Double  Decomposition  in  Solution*  —  We  are 

now  prepared  to  consider  the  general  case  of  mixing  the  solutions  of 
two  ionogens. 
When  solutions  of  two  ionized  substances  are  mixed,  the  first 


IONIC   SUBSTANCES  237 

reflection  which  occurs  to  us  is  that  each  of  these  has  been  diluted 
by  the  water  in  which  the  other  was  dissolved,  so  that  the  first  effect 
will  be  to  increase  the  degree  of  ionization  of  both  to  a  certain  extent. 
The  next  consideration  is,  however,  that  we  haveproduced  a  mixture 
of  four  ions;  which  must  have  at  least  some  tendency  to  unite  cross- 
wise. Thus  potassium  chloride  and  sodium  nitrate  in  dilute  solution 
are  very  greatly  ionized  before  mixing.  The  reversible  actions, 
represented  by  the  horizontal  pair  of  the  following  equations,  have 
taken  place  extensively.  But,  by  mixing  the  liquids,  we  have 
brought  into  presence  of  one  another  two  new  pairs  of  positive  and 
negative  ions.  Hence,  two  other  reversible  actions,  the  vertical 
ones, 


NaN03  ±^  NO/  +  Na- 

if  '         it 
KNO3      NaCl 

will  be  set  up  and  will  proceed  until  a  fresh  equilibrium  of  all  the 
ions  with  all  four  kinds  of  molecules  has  been  reached.  Thus  far 
the  description  will  fit  any  case  of  mixing  solutions  of  two  ionogens. 

Now,  in  this  particular  instance,  what  is  the  actual  extent  of  such 
interaction  as  has  occurred?  To  answer  this  question  we  require  to 
know  the  proportion  of  molecules  to  ions  in  a  solution  of  each  of  the 
four  salts.  In  decinormal  solutions  it  is  KC1,  14  :  86  ;  NaNO3,  17  :  83  ; 
KNO3,  17  :  83;  NaCl,  16  :  84,  so  that  the  salts  are  all  equally  well 
ionized.  Furthermore,  in  a  dilute  mixture,  such  as  we  are  consider- 
ing, the  proportions  of  ions  are  greater  than  these  figures  indicate. 
Hence,  practically  no  chemical  action  has  occurred. 

That  this  inference  is  correct  is  shown  by  independent  evidence. 
Thus  when  the  solutions  are  mixed,  no  thermal  effect  is  observable. 
Again,  if  the  solutions  are  placed  in  a  cell  (Fig.  53,  p.  225),  so  that 
the  one  forms  a  layer  below  the  other,  no  change  in  conductivity  is 
noticed  when  the  solutions  are  stirred  together.  Hence  no  change 
in  the  number  of  ions  has  occurred. 

We  conclude,  then,  that  when  two  highly  ionized  substances  are 
mixed,  and  the  possible  products  are  also  highly  ionized,  soluble  sub- 
stances, then  practically  no  chemical  action  occurs.  This  rule  applies 
to  all  soluble  salts  (p.  229)  and  to  the  highly  ionized  acids  and  bases. 

Conversely,  when  two  ionized  substances  are  mixed,  an  extensive 
chemical  change  does  ensue  in  two  cases,  namely: 


238  COLLEGE   CHEMISTRY 

1.  When  one  of  the  possible  products  is  an  insoluble  substance  and 
precipitation  occurs,  for  this  removes  the  ions  used  in  forming  the 
insoluble  body. 

2.  When  one  of  the  possible  products,  although  soluble,  is  little 
ionized,  as  in  neutralization,  for  this  likewise  removes  the  ions  required 
to  form  molecules  of  the  product.     We  proceed  therefore,  to  discuss 
these  two  important  classes  of  actions. 

Precipitation.  —  A  typical  case  of  precipitation  occurs  when 
we  mix  dilute  solutions  of  silver  nitrate  and  sodium  chloride: 

NaCl  ±;  Na"    +  Cl' 
AgN03  ±^  NO/  +  Ag" 

It  JI 

NaNO3        AgCl  (dslvd) 

it 

AgCl  (solid) 

Here,  since  the  four  substances  are  all  salts,  they  are  all  highly 
ionized.  If  they  were  all  soluble,  then,  in  dilute  solutions,  perhaps 
5  per  cent  of  each  salt  would  be  in  molecules  and  the  rest  in  ionic 
form.  But  the  molecules  of  silver  chloride  are  excessively  insoluble. 
In  all  cases  of  precipitation,  we  look  up  the  solubilities  of  the  possible 
products  (see  Table  of  solubilities  inside  the  front  cover).  Here  we 
find  that  one  liter  of  water  will  dissolve  only  0.0016  g.  silver  chloride. 
So  the  concentration  of  the  AgCl  (dslvd)  becomes  almost  zero  through 
precipitation.  This  displaces  the  equilibrium,  for,  the  dissociation 
having  thus  ceased,  those  of  the  ions  Ag"  and  Cl'  which  combine 
are  not  replaced  by  others.  Hence  the  silver  and  chloride-ion 
disappear.  This  occurrence  affects  in  turn  the  equilibria  with  Na" 
and  NO/,  so  that  the  NaCl  and  AgNO3  become  completely  ionized. 
Hence  the  concentrations  of  NaCl  and  AgN03,  of  Ag"  and  Cl',  and 
of  the  dissolved  AgCl,  all  become  practically  zero  at  last.  The  sys- 
tem finally  contains  only  a  precipitate  of  molecular,  solid  silver 
chloride  and  a  solution  of  the  three  substances,  Na"  +  NO/  ±3  NaNO3, 
in  equilibrium.  By  far  the  greater  part  of  this  material  in  solution 
is  the  ionic,  namely  the  Na"  and  the  NO/. 

It  should  be  noted  that,  when  the  solutions  are  mixed,  as  in  the 
foregoing  example,  strictly  speaking,  the  chief  interaction  taking 
place  is  the  production  of  the  insoluble  body.     The  largest  part  of 
the  chemical  action  may  be  formulated  thus: 
Ag"  +  Cl'  ->  AgCl. 


IONIC  SUBSTANCES  239 

The  chief  change  that  has  as  yet  befallen  the  ions  of  sodium  nitrate  is 
that  they  have  been  transferred  from  two  separate  vessels  into  one. 
Potentially  the  salt  has  been  formed.     But  the  actual  union  of  its 
ions,  to  give  the  second  product  in  the  molecular  condition: 
Na'  +  NO/  ->  NaNO3, 

comes  about  only  when,  at  some  subsequent  time,  if  at  all,  the  water 
is  evaporated  away. 

The  foregoing  formulation  and  explanation  apply  to  every  case 
of  mixing  ionogens  where  precipitation  occurs,  that  is,  where  the 
products  are  insoluble  acids,  bases,  or  salts. 

\  Neutralization.  —  We  may  now  consider  the  case  of  mixing  solu- 
tions of  two  ionogens  where  one  is  an  acid  and  one  a  base. 

The  general  plan  of  all  interactions  of  acids  and  bases  is  as  follows: 
HC1  fc;  Cl'    +  IT 

NaOH  fc?  Na'  +  OH' 

IT       IT 

NaCl     H2O 

The  ionization  of  the  hydrochloric  acid  reaches  0.91  in  a  decinormal 
solution,  and  goes  further  when  the  acid  is  diluted  with  the  water  of 
another  solution.  That  of  the  sodium  hydroxide  similarly  goes 
beyond  0.84.  Thus  the  substances  in  the  solutions  before  mixing 
are  almost  entirely  ionic.  The  crosswise  union,  H"  +  OH'  t=>  H2O, 
however,  is  all  but  complete,  for  water  is  hardly  ionized  at  all  (p.  228). 
The  materials  on  whose  interaction  with  the  Cl'  and  Na',  respectively, 
the  maintenance  of  molecules  HC1  and  NaOH  depends,  being  thus 
removed,  the  dissociation  of  the  acid  and  base  promptly  brings  itself 
to  completion,  and  the  left  sides  of  the  equations  vanish.  Practi- 
cally all  the  hydrogen-ion  and  hydroxide-ion  become  water,  which 
thenceforth  is  simply  a  part  of  the  solvent.  The  Cl'  and  Na*,  how- 
ever, if  the  solution  is  now  1/20  normal,  unite  to  the  extent  of  0.13 
only.  If  it  is  more  dilute,  this  union  forms  a  still  smaller  factor  in 
the  whole  change.  Practically  it  is  negligible.  Now  all  that  has 
been  said  of  this  acid  and  base  will  apply  mutatis  mutandis  whenever 
any  active,  highly  ionized  acid  and  base  come  together.  Thug  we 
may  write  one  simple  equation  for  all  neutralizations  of  active  acids  and 

ba*es:  H'  +  OH'-»H20, 

without  omitting  anything  essential. 


240  COLLEGE   CHEMISTRY 

The  ions  of  a  salt  are  always  left  over  from  the  main  action,  and 
may  be  brought  together,  in  turn,  by  evaporation :  Na"  +  Cl'  — >  NaCl, 
or  the  liquid  may  be  used  as  a  solution  of  the  pure  salt. 

That  these  inferences  are  correct  is  shown  by  many  facts.  The 
most  conspicuous  of  these  is  the  fact  that,  when  equivalent  amounts 
of  the  acids  and  bases  are  used,  the  mixture  is  without  action  either 
on  red  or  on  blue  litmus.  It  is  neutral  to  indicators  —  hence  the  term 
neutralization  applied  to  the  operation  of  mixing  an  acid  and  a  base. 
Specifically,  the  absence  of  effect  upon  litmus  demonstrates  the 
absence  of  hydrogen-ion  H*  and  of  hydroxide-ion  OH',  alike,  in  the 
product,  and  confirms  the  theory. 

Again,  a  considerable  thermal  effect  accompanies  neutralization. 
But,  in  the  cases  we  are  discussing,  that  is  where  active  bases  and 
acids  are  employed,  the  heat  liberated  by  use  of  equivalent  weights 
(p.  99)  is  always  the  same,  namely  13,700  cal.  That  it  is  always 
the  same  confirms  our  theory,  for  practically  the  whole  change  is 
always  the  formation  of  18  g.  of  water  from  the  ions. 

Still  again,  when  we  place  the  acid  and  base  in  the  cell  (Fig.  53, 
p.  225),  so  that  the  one  forms  a  layer  beneath  the  other,  and  watch 
the  amperemeter  while  we  mix  the  solutions,  a  marked  decrease  in 
the  current  passing  through  the  cell  is  noticed.  This  also  confirms 
our  theory,  for  it  is  our  belief  that  one-half  of  the  ions,  namely  the 
H*  and  OH',  disappear  as  such  during  the  action. 

When  less  highly  ionized  acids  or  bases  are  used,  the  only  difference 
is  that  there  are  more  of  the  molecular  materials  present,  before  the 
solutions  are  mixed.  But  the  removal  of  the  H*  and  OH'  ions  per- 
mits the  molecules  of  the  acid  and  base  to  dissociate,  so  that  the 
final  products  are  water  and  the  ions  of  a  salt,  as  before. 

The  foregoing  formulation  and  explanation  apply  to  every  case  of 
mixing  ionogens,  where  a  very  slightly  ionized  substance  is  one  of 
the  products,  that  is,  when  water,  or  a  feeble  acid,  or  a  feeble  base 
(pp.  229-230)  is  formed. 

Acidimetry  and  Alkalimetry.  —  When,  as  is  constantly  the 
case,  a  chemist  desires  to  ascertain  the  quantity  of  an  acid  or  base 
present  in  a  solution,  he  uses  for  the  purpose  the  interaction  just 
discussed.  If,  for  example,  the  problem  is  to  ascertain  the  weight 
of  hydrogen  chloride  in  each  liter  of  a  specimen  of  hydrochloric  acid, 
this  can  be  done  by  neutralizing  a  measured  portion  of  this  acid  with 


IONIC    SUBSTANCES 


a  solution  of  an  alkali  of  known  concentration.  The  volume  of  the 
latter  which  is  required  for  the  purpose  is  observed.  If  the  alkali  is 
sodium  hydroxide,  the  action  taking  place  is: 

HC1  +  NaOH  ->  H2O  +  NaCl. 

The  volume  of  acid  is  measured  out  into  a  beaker  by  means  of  a 
pipette  (Fig.  54)  of  fixed  capacity,  which  is  filled  by  suction  to  the 

mark  on  the  stem.     Suppose  the  amount  to  be  25  c.c.     The 

standard  alkali  solution  is  placed  in  a  burette  (Fig.  55), 

which  is  filled  down  to  the  tip  of  the  nozzle.     A  few  drops 

of  litmus  solution  are  now 

added  to  the  acid,  and  the 

alkali  is  allowed  to  run  in 

slowly.     After  a  time,  the 

hydroxide-ion  which  this 

introduces   will   begin  to 

produce  a  blue  color,  close 

to  where  the  stream  enters 

the  liquid.     This  is  at  first 

dissipated  by  stirring,  and 

the  whole  remains    red. 

Finally,  however,  a  point 

is  reached  at  which  the 

entire  solution  assumes  a 

tint  intermediate  between 

blue  and  red.     With  one 

drop  less  of  the  base,  it  is 

distinctly  red.     With  one 

drop  more,  it  would  be- 
come distinctly  blue.  Lit- 
mus paper  of  either  shade 

dipped  in  this  neutral  solu- 
tion remains  unaffected. 
FIG.  54.       By  the  use  of  a  standard 

solution  of  an  acid,  the 
quantity  of  a  base  may  be  deter-  FlQ- 55- 

mined  in  the  same  way. 

The  standard  solutions  used  in  this  work  are  usually  normal,  and 
contain  one  equivalent  weight  of  the  alkali  or  acid  in  one  liter  of  the 


242  COLLEGE   CHEMISTRY 

solution.  For  more  delicate  work,  decinormal  solutions  may  be 
employed.  The  concentration  of  such  a  solution  is  called  its  titer, 
and  the  operation  of  analyzing  another  solution  by  means  of  it, 
titration.  The  value  of  standard  solutions  lies  in  the  fact  that,  when 
once  the  solution  has  been  prepared,  and  the  exact  concentration 
adjusted  by  quantitative  experiments,  its  use  does  not  require  any 
weighing,  and  the  measurements  of  volumes  can  be  carried  out  with 
great  rapidity.  The  calculation  of  the  result  is  also  simple.  One  liter 
of  normal  alkali  contains  17  g.  of  available  hydroxyl,  and  one  liter  of 
normal  acid,  1  g.  of  available  hydrogen  (p.  99).  Equal  volumes 
of  normal  solutions  will  therefore  exactly  neutralize  one  another, 
18  g.  of  water  being  formed  by  interaction  of  a  liter  of  each.  If,  for 
the  neutralization  of  the  25  c.c.  of  hydrochloric  acid  used  above,  50 
c.c.  of  normal  alkali  are  required,  the  acid  is  twice-normal  (2N). 
When  15  c.c.  are  required,  the  acid  is  £f  or  f  N.  If  the  actual  weight 
of  the  acid  in  the  latter  case  has  to  be  calculated,  we  remember  that 
there  are  36.45  g.  of  hydrogen  chloride  in  1  1.  of  a  normal  solution, 
and  therefore  36.45  X  f  X  T§jhf  g-  =  -5467  g.  in  25  c.c.  of  a  solution 
which  is  f-normal. 

Methods  of  quantitative  analysis  in  which  standard  solutions  are 
employed  are  known  as  volumetric  methods,  and  are  much  used  by 
analysts  and  investigators.  They  occupy  much  less  time  than  gravi- 
metric operations,  in  which  numerous  weighings  have  to  be  made, 
and  are  often  just  as  accurate.  The  substances,  like  litmus,  by  whose 
change  of  color  the  completeness  of  the  action  is  made  known,  are 
called  indicators  (see  below). 

Indicators.  —  Indicators  are  substances  which,  in  presence  of 
certain  other  substances,  assume  a  very  deep  color,  or  change  sharply 
from  one  deep  color  to  another.  Thus,  phenolphthale'in  is  colorless 
in  presence  of  acids  (i.e.,  hydrogen-ion),  and  red  (when  dilute,  pink) 
in  presence  of  alkalies  (i.e.,  hydroxide-ion).  Litmus,  again,  is  red 
with  acids,  and  blue  with  alkalies.  The  change  of  color  depends 
upon  a  chemical  interaction  in  each  case,  but  since  indicators  are 
chosen  for  their  strong  coloration,  the  quantity  of  the  acids  or  base 
used  up  in  changing  the  tint  of  the  trace  of  the  indicator  is  so  small 
as  to  be  negligible.  The  common  indicators  are: 

Phenolphthaleih,  C14H10O4,  a  colorless  substance  and  very  feeble 
acid.  It  is  not  perceptibly  dissociated  into  its  ions: 

C14H10O4  (colorless)  ±=>C14H9O4'  (red)  +  H', 


IONIC   SUBSTANCES  243 

and  in  neutral  or  acid  solutions  is,  therefore,  without  visible  color. 
When  a  base  is  added  gradually  to  an  acid  containing  some  of  this 
indicator,  the  acid  is  first  neutralized.  Then,  and  not  till  then,  the 
slightest  excess  of  hydroxide-ion  unites  with  the  trace  of  hydrogen- 
ion  from  the  phenolphthalem,  the  above  equilibrium  is  displaced 
forwards,  and  a  visible  amount  of  the  red  negative  ion  is  formed: 

C14H1004  (colorless)  ^C14H90/ (red)  +   ET   j  <-HO 
NaOH  *=»Na'  +  OH'J  -"W^ 

Litmus  is  a  natural  dyestuff  of  unknown  chemical  structure.  One 
of  its  colors  is  that  of  the  molecule,  and  the  other  that  of  the  ion. 

Methyl  orange,  (CH3)2NC6H4.N  :  N.C6H4SO3Na,  is  a  complex 
organic  compound  which  gives,  in  acid  solution,  a  red,  and  in  alkaline 
solution  a  yellow  color. 

Congo  red  is  the  sodium  salt  of  an  acid  of  complex  structuje  (see 
Dyes).  In  neutral  or  alkaline  solutions  it  is  red;  with  acidsit  turns 
blue.  Paper  dipped  in  Congo  red  differs  from  litmus  paper  in  that  it 
shows  gradations  in  color,  the  blue  being  much  more  distinct  with  an 
active  acid  than  with  a  relatively  weak  one  like  acetic  acid  (p.  229). 
Litmus  paper  is  equally  red  with  all  acids  save  the  very  feeblest. 

Displacement :  The  Electromotive  Series. —  In  the  preced- 
ing sections  we  have  dealt  with  cases  in  which  ionic  substances  under- 
went combination  or  ionogens  dissociated.  This  is  one  of  five  kinds  of 
ionic  chemical  change.  Of  the  remaining  four,  ionic  displacement  is 
the  one*  that  we  have  most  frequently  encountered.  Thus,  certain 
metals  displace  hydrogen  from  dilute  acids  (p.  66) : 
Zn  +  H2SO4  -»  ZnS04  +  H2. 

These  interactions  do  not  occur  in  the  absence  of  water  (p.  65),  and 
now  appear  in  a  new  light,  namely,  as  ionic  actions: 

Zn  +  2IT  +  SO/'  ->  Zn"  +  H2  +  S04". 

The  molecular  sulphuric  acid  and  zinc  sulphate,  which  are  small  in 
amount,  are  omitted  because  they  do  not  take  part  in  the  change. 
On  looking  at  the  equation,  we  perceive  that  the  sulphate-ion  is  also 
unaltered  by  the  action,  and  may  be  left  out  likewise: 
Zn  +  2H'  -»  Zn"  +  H2. 

*  The  discharge  of  an  ion  and  liberation  of  its  material  in  electrolysis 
(pp.  64,  109,  215)  is  another.  Attention  will  be  called  to  the  remaining  two 
when  suitable  illustrations  occur  (see  pp.  253,  412). 


244  COLLEGE    CHEMISTRY 

True,  hydrogen-ion  cannot  be  used  alone,  for  it  is  always  accom- 
panied by  some  negative  radical.  But  the  latter,  like  the  vessel  in 
which  the  experiment  is  made,  is  part  of  the  necessary  apparatus, 
and  not  an  interacting  substance.  The  change  has  consisted  in  the 
ionization  of  the  zinc,  and  the  transfer  to  it  of  the  electric  charge  of 
the  hydrogen-ion. 

These  statements  enable  us  to  understand  why  active  acids,  with 
zinc,  give  hydrogen  faster  than  do  inactive  acids  (p.  65).  The 
former  provide-  a  higher  concentration  (p.  229)  of  hydrogen-ion, 
that  is,  of  the  real  interacting  substance,  than  do  the  latter. 

A  similar  displacement  of  negative  ions  has  been  met  with  (pp.  159, 
164).  Thus,  chlorine  displaces  bromine  from  solutions  containing 
bromide-ion.  CL  +  2Br- -*  2C1' +  Br, 

Displacement  occurs  with  all  positive  ions.  Thus,  zinc  will  dis- 
place other  metallic  elements,  such  as  iron,  lead,  copper,  and  silver, 
from  the  ionic  conditions,  when  it  is  placed  in  solutions  of  their  salts: 

Zn  +  Cu'*  -»  Zn"  +  Cu. 

Here  the  copper  appears  as  a  red  precipitate.  Lead,  in  turn,  will 
displace  copper  and  silver,  but  not  zinc  or  iron.  Copper  will  displace 
silver.  Thus  the  metals  can  be  set  down  in  an  order,  such  that  each 
metal  displaces  those  following  ijb  in  the  list  and  is  displaced  by  those 
preceding  it.  This  list  (see  next  page)  is  known  as  the  electromotive 
series  of  the  metals,  because  in  electrolysis  of  normal  solutions  of 
their  salts,  the  electromotive  force  of  the  current  required  to  deposit 
each  metal  is  less  than  that  for  the  metal  preceding  in  the  list.  For 
present  purposes,  the  list  shows  the  metals  in  the  order  of  diminishing 
tendency  to  enter  the  ionic  from  the  elementary  condition. 

The  electromotive  series  embodies,  many  facts  in  the  behavior  of 
the  metals,  and  should  be  kept  in  mind  as  furnishing  a  key  to  all 
actions  in  which  a  fre'e  metal  is  used  or  produced.  For  example,  the 
chemical  activity  of  the  free  metals  places  them  in  the  same  order. 
The  earliest  ones  rnst  much  more  readily  in  air  ftian  do  the- later  ones. 
Those  following  copper  do  not  rust.  Conversely,  the  oxides  of  the 
metals  down  to  and  including  manganese',  when  heated  in  a  stream 
of  hydrogen,  may  give  lower  oxides,  but  are  not  completely  reduced. 
The  oxides  of  cadmium  .and  succeeding  metals  .are  easily  reduced. 
The  oxides  of  mercury,  and  the  last  ^our  metals  are  decomposed 


IONIC    SUBSTANCES 


245 


by  heating  alone.     The  relations  of  the  metals  in  respect  to  com- 
bination with  elements  other  than  oxygen  are  similarly  expressed 
by  the  arrangement  in  this  table. 

The  position  of  hydrogen  is  particularly  sig- 
nificant. It  will  be  noted  that  none  of  the  metals 
preceding  hydrogen  are  found  free  in  nature  as 
ordinary  minerals,  while  all  of  the  metals  suc- 
ceeding hydrogen,  although  occurring  to  some 
extent  in  combination,  are  found  also  free.  The 
explanation  of  this  is  that,  by  prolonged  action 
upon  ordinary  water,  containing,  as  it  must,  car- 
bonic acid  and  other  sources  of  hydrogen  ions, 
the  metals  preceding  hydrogen  must  eventually 
displace  hydrogen-ion  and  pass  into  some  form  of 
combination.  The  metals  following  hydrogen 
do  not  displace  hydrogen-ion  and  are  therefore 
much  less  affected  by  the  agencies  which  are  most 
active  in  the  chemical  transformation  of  min- 
erals. Hence  they  often  remain  in  the  free 
state. 

The  negative  ions  can  be  arranged  in  order  in 
a  similar  way. 

To  avoid  a  common  misconception,  it  must 
be  noted  that  the  electromotive  series  cannot 
be  used  to  explain  the  tendency  of  one  radical 
to  dislodge  another  in  double  decompositions. 
The  place  of  an  element  in  the  E.M.  series  defines 
its  relative  activity  when  free,  and  has  to  do  only  with  actions 
where  one  free  Clement  displaces  (p.  68)  another.  The  influences 
which  determine  a  double  decomposition  (c/.  pp.  184,  187)  are  such 
as  the  insolubility  of  a  compound.  Thus,  potassium  bromide  solution 
will  slowly  convert  a  precipitate  of  silver  chloride  into  one  of  silver 
bromide:  AgCl-}-KBr-»AgBr'+KCl.  Th'is  occurs  because  silver 
bromide  is  the  less  soluble  salt.  But  free  bromine  never,  displaces 
chlorine  from  binary  combination  with  a  metallic  element. 

"rf 

Non-Ionic  Modes  of  Forming  lonogens. — .While  ionogens 
may  always  be  made  by  the  union  of  the  proper  ions,  they  must 
nevertheless,  in  the  absence  of  the  solvent,  be  regarded  as  chemical 


ELECTROMOTIVE 
SERIES  OF  THE 
METALS. 

Alkali  metals  (q.v.) 
Alkaline  earth 
metals  (q.v.) 
Magnesium 
Aluminium 
Manganese 
Zinc 

Chromium 
Cadmium 
Iron 
Cobalt 
Nickel 
Tin 
Lead 

Hydrogen 
Arsenic 
Copper 
Antimony 
Bismuth 
Mercury 
Silver  — 
Palladium 
Platinum 
Gold 


246  COLLEGE  CHEMISTRY 

substances  which  may  be  constructed  out  of  their  constituents  with- 
out reference  to  the  ionic  plane  of  cleavage.  Thus  we  have  inciden- 
tally observed  many  ways  in  which  acids,  bases,  and  salts  may  be 
prepared,  that  do  not  involve  a  union  of  the  constituent  ions  and  are 
probably  not  ionic. 

Oxygen  acids  can  almost  all  be  prepared  from  the  anhydrides, 
which  are  not  ionogens,  and  water.  Phosphoric  acid,  sulphurous 
acid  (p.  51),  hypochlorous  acid  (p.  191),  and  many  other  acids  are 
so  formed.  Hydrogen  fluoride,  chloride,  bromide,  and  iodide  are 
all  producible  by  union  of  the  constituent  elements.  Many  acids 
are  formed  from  others  when  the  latter  are  heated;  for  example, 
perchloric  acid  from  chloric  acid  (p.  196). 

Bases  are  formed  by  the  union  of  oxides  of  metals  with  water 
(p.  81). 

The  dry  ways  of  forming  salts  are  very  numerous.  Thus,  many 
are  produced  by  direct  union  of  the  elements,  as  in  the  case  of  chlo- 
rides (p.  113),  sulphides  (p.  29),  and  other  simple  salts.  Many  are 
made  by  reduction  or  oxidation  from  other  salts,  as  potassium  chlo- 
ride from  potassium  chlorate  (p.  47),  or  potassium  perchlorate  from 
the  latter  (p.  196).  Often  a  reducing  or  an  oxidizing  agent  is  used, 
as  in  making  sodium  nitrite  (q.v.)  from  the  nitrate.  Almost  all 
oxygen  salts  can  be  obtained  by  the  union  of  two  oxides,  as  calcium 
carbonate  (q.v.)  from  calcium  oxide  and  carbon  dioxide.  Ammo- 
nium salts  are  formed  by  combination  of  ammonia,  which  is  not  an 
ionogen,  with  acids  (p.  121). 

In  manufacturing  salts,  methods  like  the  above,  as  well  as  those 
involving  ionic  actions,  are  very  commonly  used.  In  each  case  the 
cheapest  and  most  easily  accessible  materials  are  chosen,  and  the 
least  expensive  operation  is  selected. 

Exercises.  —  1 .  Give,  for  each  of  the  following,  a  definition, 
i.e.  concise  description,  in  terms  of  experimental  facts:  acid  (pp.  64, 
111,  187,  201),  base  (pp.  81,  201),  salt  (p.  187),  acid  salt,  mixed  salt. 

2.  Give,  now,  a  definition  of  the  same  things  (see  1),  in  terms  of 
the  hypothesis  of  ions. 

3.  Name  all  the  ionic  substances  whose  formulae  are  given  on 
pp.  203,  223,  and  classify  them  into  anions  and  cations. 

4.  Give  a  list  of  the  specific  physical  and  chemical  properties 
including  those  that  can  be  used  as  tests,  of:  iodide-ion,  sulphate- 
ion,  cupric-ion,  peroxide-ion. 


IONIC  SUBSTANCES  247 

5.  Give  a  list  of  all  the  colorless  ionic  substances  you  can  think  of. 

6.  Using  the  table  of  fractions  ionized  (p.  228),  prepare  lists  of 
the  pairs  of  ionic  substances  which  show  the  greatest,  and  the  least 
tendency  to  combine,  and  state  in  each  case  the  proportion  com- 
bining in  decinormal  solution. 

7.  In  the  case  of  the  green  solution  of  cupric  bromide  (p.  235), 
explain  in  detail  (p.  179)  the  effect  of  the  addition  of  potassium 
bromide. 

8.  In  the  case  of  the  chocolate-brown,  concentrated  solution  of 
cupric  bromide  (p.  235),  explain  in  detail  what  would  happen  to  the 
system:  (a)  if  metallic  zinc  were  to  be  added  (p.  244);  (b)  if  hydrogen 
sulphide  gas  were  to  be  led  into  the  solution  (CuS  is  insoluble). 

9.  Formulate,  after  the  models  on  pp.  237  and  238,  and  discuss 
fully:  (a)  the  interaction  of  dilute  sulphuric  acid  and  potassium 
permanganate  (p.  213);  (b)  the  preparation  of  chloric  acid  (p.  195). 

10.  What  is  implied  by  the  statements,  that  peroxides  are  salts 
and  that  hydrogen  peroxide  is  feebly  acid  (p.  212)? 

11.  Formulate  after  the  model  on  p.  238,  and  discuss  fully,  the 
interaction  of:   (a)  sodium  peroxide  and  hydrochloric  acid  (p.  211); 
(b)  barium  peroxide  and  sulphuric  acid. 

12.  Can  you  invent  an  interaction  of  two  soluble  salts  in  which 
both  products  shall  be  insoluble  (see  Table  of  solubilities,  inside  of 
front  cover)? 

13.  For  the  neutralization  of  77  c.c.  of  a  certain  alkaline  solution, 
25  c.c.  of  normal  hydrochloric  acid  are  required.     What  is  the  normal 
concentration  of  the  alkali?    If  the  alkali  was  sodium  hydroxide, 
what  weight  of  the  substance  was  present?    If  the  alkali  was  barium 
hydroxide,  what  weight  of  it  was  present? 

14.  Formulate  (p.  243)  the  actions  of  iron  and  of  aluminium  on 
dilute  hydrochloric  acid. 

15.  Formulate  (p.  244)  the  displacements  of  iodine  by  chlorine 
and  by  bromine  (p.  166). 

16.  Which  metals  (p.  245),  beside  platinum,  would  be  most  likely 
to  form,  suitable  electrodes  for  an  electrolytic  cell? 

17.  To  which  classes  of  ionic  actions  do  those  of  iodine  on  hydro- 
gen sulphide  (p.  167),  and  of   magnesium  on  cold   water  (p.  66), 
belong? 

18.  Give  all  the  different  ionic  interactions  by  which,  (a)  acids, 
(b)  bases,  and  (c)  salts  are  prepared,  with  illustrations  of  each. 


\r 


CHAPTER  XXI 
.•• 
SULPHUR   AND    HYDROGEN    SULPHIDE 


Occurrence.  —  Free  sulphur  is  found  in  volcanic  regions,  where  it 
is  mixed  with  gypsum  and  other  minerals  and  occupies  the  pores  of 
pumice-stone.  Rocky  materials  accompanying  a  mineral  in  this  way 
are  called  the  matrix.  Here  and  there  non-volcanic  deposits,  formed 
by  the  action  of  bacteria,  have  been  met  with,  as  in  Louisiana  and  in 
Germany.  There  are  many  minerals,  compounds  containing  sul- 
phur, which  are  chiefly  important,  however,  on  account  of  their  other 
constituents.  Sulphides  of  metals,  such  as  pyrite  (FeS2),  copper 
pyrites  (CuFeS2),  galena  (PbS),  zinc-blende  (ZnS),  and  sulphates, 
like  gypsum  (CaSO4,2H2O),  barite  (BaSO4),  and  celestite  (SrSO4), 
are  fairly  plentiful.  Sulphur  is  a  constituent  of  albumin  and  other 
substances  found  in  the  animal  body. 

Manufacture.  —  Most  sulphur  is  obtained  by  the  simple  process 
of  melting  it  away  from  the  accompanying  volcanic  rock  at  a  low 
temperature.  The  liquid  sulphur  is  allowed  to  run  into  wooden 
molds,  in  which  it  solidifies  in  the  form  of  roll  sulphur.  To  produce 
the  best  quality  it  is  subjected  to  distillation  from  earthenware 
retorts.  When  the  vapor  is  led  into  a  large  brick  chamber,  it  con- 
denses upon  the  walls  and  floor  in  the  form  of  flowers  of  sulphur. 

The  greater  part  of  the  sulphur  of  commerce  comes  from  Sicily, 
where,  in  1898,  447,000  tons  were  manufactured  against  41,000  tons 
elsewhere.  It  is  found  in  Japan,  California,  Nevada,  and  Louisiana. 
Sulphur  is  popularly  known  as  brimstone. 

Physical  Properties.  —  The  chief  physical  peculiarity  of  sulphur 
is  that,  instead  of  existing  in  three  physical  states  only,  like  water,  it 
possesses  two  familiar  and  perfectly  distinct  solid  forms  and  two 
different  liquid  states  of  aggregation. 

1.  Native  sulphur  is  yellow,  has  a  sp.  gr.  2.06  and  melts  at  114.5°. 
It  is  almost  insoluble  in  water,  but  dissolves  freely  in  carbon  disul- 

248 


SULPHUR  249 

phide  (41  parts  in  100  at  18°).  The  crystals  of  native  sulphur,  as 
well  as  those  obtained  by  evaporating  a  solution,  belong  to  the  rhom- 
bic system  (Fig.  1,  cf.  p.  6).  Roll  sulphur  is  the  same  substance 
as  these  two,  although  the  crystals  in  their  growth  have  interfered 
with  one  another,  and  the  mass  is  crystalline,  simply,  and  not  well 
crystallized.  This  variety  is  called,  fron\its  form,  rhombic  sulphur. 

2.  When  a  large  mass  of  melted  sulphur' solidifies  slowly,  and  the 
crust  is  pierced  and  the  remaining  liquid  pourebV  out  before  the  whole 
has  become  solid,  the  interior  is  found  to  be  lined  with  long,  trans- 
parent needles.     This  kind  of  sulphur  is  nearly  colorless  and  has  a 
sp.  gr.  1.96,  melts  at  119°,  and  is  in  all  physical  respects  a  different 
individual  from  rhombic  sulphur.     This  variety  is  named,  from  the 
system  to  which  its  crystals  belong,  monoclinic  sulphur. 

A  substance  which  has  two  solid  states  of  aggregation  and,  there- 
fore, two  crystalline  forms,  is  said  to  be  dimorphous  (two-formed). 

3.  When  melted  sulphur  is  heated,  it  undergoes  another  change, 
which  is  especially  noticeable  near  160°.     The  formerly  pale-yellow, 
mobile  liquid  (S*)  suddenly  becomes  dark-brown  in  color  and  so  vis- 
cous (Sfj,)  that  the  vessel  may  be  inverted  without  loss  of  material. 
Beyond  260°  the  viscidity  becomes  less,  and  at  445°  the  liquid  boils 
and  passes  into  sulphur  vapor. 

When  ordinary  sulphur  is  boiled  and  then  allowed  slowly  to  cool, 
the  product  is  crystalline  and  soluble  in  carbon  disulphide,  as  before. 
But  when  sulphur  is  boiled  and  then  suddenly  chilled  by  pouring  into 
cold  water,  it  is  at  first  semi-fluid.  After  several  days  this  elastic 
sulphur,  as  it  is  called,  becomes  hard.  It  is  then  found  to  contain 
rhombic  sulphur  mixed  with  a  large  proportion  of  another  variety 
of  free  sulphur.  This  is  almost  insoluble  in  any  solvent.  Being 
without  crystalline  structure,  it  is  called  amorphous  (Gk.  d  priv., 
l*op<j>rj  form)  sulphur.  Now  amorphous  bodies  (see  Glass)  are  always 
supercooled  liquids,  that  is,  liquids  still  existing  as  such  at  a  tem- 
perature at  which  the  solid,  crystalline  form  is  the  stable  one.  This 
is  simply  the  brown,  viscous  sulphur  (SM)  in  a  supercooled  state.  It 
reverts  very  slowly  to  the  soluble  variety,  and  years  are  required  for 
the  completion  of  the  reversion  at  room  temperature. 

Chemical  Properties.  —  At  low  temperatures  and  under  reduced 
pressure,  the  formula  of  sulphur  vapor  is  S8.  As  the  temperature  is 
raised,  however,  the  vapor  expands  very  rapidly,  and  at  800°  the 


\ 


250  COLLEGE   CHEMISTRY 

molecular  weight  is  64.2,  and  the  formula  therefore  S2  (p.  146). 
The  formula  of  dissolved  sulphur,  as  measured  by  the  freezing-point 
method  (p.  205),  is  Sg. 

Sulphur  is  an  active  chemical  substance.  When  finely  divided 
metals,  with  the  exception  of  gold  and  platinum  (cf.  p.  245),  are 
rubbed  together  with  powdered  sulphur,  union  takes  place  and  sul- 
phides are  produced.  Sulphur  when  heated  combines  with  great 
vigor  with  iron  (p.  6),  copper,  and  most  of  the  metals.  It  unites 
also  with  many  of  the  non-metals.  Thus  with  oxygen  it  produces 
sulphur  dioxide  (p.  48),  and  even  sulphur  trioxide  SO3.  It  unites 
also  with  chlorine  directly.  When  sulphur  is  treated  with  oxidizing 
agents  in  presence  of  water,  no  trace  of  sulphur  dioxide  (or  sulphurous 
acid)  is  formed;  the  only  product  is  sulphuric  acid.* 

Uses  of  Sulphur.  —  Large  quantities  of  crude  sulphur  are 
employed  for  making  sulphur  dioxide,  which  is  used  in  the  manu- 
facture of  sulphuric  acid,  in  bleaching  feathers,  straw,  and  wool,  and 
in  making  alkali  sulphites  for  employment  in  the  bleaching  industry. 
The  manufacture  of  carbon  disulphide  consumes  a  considerable 
amount  also.  The  purified  sulphur  is  employed  in  the  manufacture 
of  gunpowder,  fireworks,  matches,  and,  by  combination  with  rubber, 
of  vulcanite.  Flowers  of  sulphur  is  used  in  vineyards  to  destroy 
fungi,  which  it  does  by  virtue  of  the  traces  of  sulphuric  acid  it  yields 
by  oxidation. 

HYDROGEN  SULPHIDE  H2S. 

This  gas  is  found  dissolved  in  some  mineral  waters,  which  in  con- 
sequence are  known  as  sulphur  waters.  It  is  produced  in  the  decom- 
position of  animal  matter  containing  sulphur,  when  air  is  excluded. 
Hence  the  distinctive  odor  of  rotten  eggs  is  due  in  part  to  its  presence. 


Preparation.  —  1.  Hydrogen  and  sulphur  do  not  unite  percep- 
tibly in  the  cold.  At  310°  almost  complete  union  occurs,  but  about 
168  hours  are  required  for  the  attainment  of  equilibrium. 

2.  Sulphides  of  metals,  being  salts,  are  acted  upon  more  or  less 
easily  by  dilute  acids  (p.  187),  and  give  hydrogen  sulphide.  Ferrous 
sulphide,  the  least  expensive  of  those  easily  affected,  is  generally 
used:  FeS  +  2HC1±;  H2S|  +  FeCl2. 

*  The  paragraph  on  the  chemical  relations  of  the  element  (see  end  of  this 
chapter)  should  be  read  at  this  point. 


HYDROGEN   SULPHIDE  251 

For  hydrochloric  acid  we  may  substitute  an  aqueous  solution  of  any 
active,  non-oxidizing  acid.  A  Kipp's  apparatus  (p.  65)  is  com- 
monly employed.  Since  ferrous  sulphide  is  but  slightly  soluble  in 
water,  the  action  proceeds  by  a  rather  complex  series  of  equilibria: 

FeS  (solid)  <=»  FeS  (dslvd)  <=»  Fe"  +  S"  L    rr  q  M^A\  _*  pr  <a  /      >> 
2HC1  *=>  2C1'  +  2H' )  ^  H2S     slvd)  ^  H2S  (gas)* 

The  dissolved  hydrogen  sulphide  is  very  feebly  ionized,  and  main- 
tains a  smaller  concentration  of  sulphide-ion  S"  than  does  ferrous 
sulphide,  in  spite  of  the  comparative  insolubility  of  the  latter. 
Hence,  the  S"  formed  from  the  FeS  is  continuously  removed  by 
union  with  the  hydrogen-ion  furnished  by  the  acid,  S"  +  2H"  t^  H2S, 
and  all  the  other  equilibria  are  constantly  displaced  forwards  on 
this  account.  The  action  is  therefore,  in  essence,  like  neutralization 
(p.  239). 

3.  Hydrogen  sulphide  is  the  invariable  product  of  the  extreme 
reduction  of  any  sulphur  compound.  Thus,  it  is  formed  by  the 
action  of  hydrogen  iodide  upon  concentrated  sulphuric  acid  (p.  166). 
Even  sulphur  itself  is  reduced  by  dry,  gaseous  hydrogen  iodide: 

2HI  +  S  ->  H2S  +  I2. 

Physical  Properties.  —  Hydrogen  sulphide  is  a  colorless  gas 
with  a  characteristic  odor.  When  liquefied,  it  boils  at  —  60.4°  (755 
mm.),  and  in  solid  form  melts  at  —  82.9°.  The  solubility  in  water 
at  10°  is  360  volumes  in  100,  and  becomes  less  as  the  temperature  is 
raised.  The  gas  can  be  driven  out  completely  by  boiling  the  solution 
(cf.  p.  120).  The  gas  is  very  poisonous,  one  part  in  two  hundred  of 
air  being  fatal  to  mammals. 

Chemical  Properties  of  Hydrogen  Sulphide  Gas.  -  -  When 
'heated,  the  gas  dissociates: 

H2S  *±  H2  +  S. 

At  310°  the  decomposition  is  slight  (cf.  p.  81),  but  becomes  greater 
at  higher  temperatures. 

The  gas  burns  in  air,  forming  steam  and  sulphur  dioxide.  The 
temperature  of  the  mantle  of  flame  surrounding  the  gas,  as  it  issues 
from  a  jet,  being  far  above  310°,  the  gas  in  the  interior  is  dissociated 
before  it  meets  with  any  oxygen.  Hence  a  cold  dish  held  across  the 


252  COLLEGE    CHEMISTRY 

flame  receives  a  deposit  of  free  sulphur,  and  a  part  of  the  hydrogen 
also  escapes  unburnt.  It  may  be  remarked  that  dissociation  of  this 
kind  probably  precedes  the  combustion  of  most  gaseous  compounds 
(see  Flame). 

The  metals,  down  to  and  including  silver  in  the  electromotive 
series,  when  exposed  to  the  gas,  quickly  receive  a  coating  of  sulphide. 
That  the  gas  should  thus  behave  like  free  sulphur  shows  its  insta- 
bility. 

This  instability  is  shown  also  in  the  fact  that  it  reduces  substances, 
such  as  sulphur  dioxide,  which  are  not  affected  by  free  hydrogen: 

2H2S  +  S02  ->  2H2O  +  3S. 

This  action  takes  place  in  the  cold,  and  much  more  rapidly  when  the 
gases  are  moist  than  when  they  are  dry  (p.  114).  Native  sulphur  is 
probably  produced  by  this  action,  as  both  of  these  gases  are  found 
issuing  from  the  ground  in  volcanic  neighborhoods.  Sulphur  is 
deposited  also  when  hydrogen  sulphide  undergoes  a  partial  combus- 
tion with  a  restricted  supply  of  oxygen,  2H2S  +  02  — >  2H20  +  2S, 
and  its  formation  in  nature  is  sometimes  to  be  accounted  for  in  this 
way. 

Chemical  Properties  of  the  Aqueous  Solution  of  Hydrogen 
Sulphide.  —  While  the  gas  itself  is  not  an  acid,  its  solution  in  water 
gives  a  feeble  acid  reaction  with  litmus,  and  is  sometimes  named 
sulphydric  acid  H2S,  Aq.  In  a  N/10  aqueous  solution,  only  .0007 
(0.07  per  cent)  of  the  substance  is  ionized: 

H2S  <=>  H*  +  HS'  (<=±  H'  +  S"). 

Some  S"  ions  are  present.  But  hydrosulphide-ion  HS',  although 
an  acid,  is  less  dissociated  than  is  water  itself,  and  the  amount  of 
sulphide-ion  is  therefore  very  small.  The  salts  of  hydrosulphide-ion, 
such  as  NaHS  (sodium  acid  sulphide,  see  next  section),  give  therefore 
neutral  solutions.  This  behavior  is  the  rule  with  the  acid  salts 
of  feeble  dibasic  acids  (p.  231). 

As  an  acid,  the  solution  of  hydrogen  sulphide  may  be  neutralized 
by  bases.  For  the  same  reason  it  enters  into  double  decomposition 
with  salts  (see  next  section). 

By  the  action  of  oxygen  from  the  air  upon  an  aqueous  solution  of 
hydrogen  sulphide,  the  sulphur  is  slowly  displaced  and  appears  in  the 
form  of  a  fine  white  powder: 

O2-i-2H,S->2SJ 


I  HYDROGEN    SULPHIDE  '    253 

This  is  an  action  similar  to  the  displacement  of  ionic  bromine  by  free 

chlorine  (p.  244). 

The  solution  of  the  gas  is  a  reducing  agent,  as  its  action  upon 

iodine  shows  (p.  167).     So,  also,  in  presence  of  an  acid,  it  removes 

oxygen  from  dichromic  acid  (produced  by  the  action  of  an  acid  upon 

potassium  dichromate) : 

KaCraO7  +  2HC1  +±  H2Cr2O7  +  2KC1  (1) 

H2Cr2O7  +  6HC1  ->  4H2O  +  2CfCl3(+  3O)         (2) 

\  /  (3O)       +  3H2S  ->  3H2O  +  3S (3) 

AddingV  K2Cr2O7  +  8HCi  +  3H2S  -*  2KC1  +  2CrCl3  +  7H2O  +  3S. 

The  first  partial  equation  (cf.  p.  159)  represents  the  regular  inter- 
action of  two  ionogens,  but  the  second  interaction  does  not  take 
place  unless  an  oxidizable  body  (here  the  hydrogen  sulphide)  is 
present  to  take  possession  of  the  oxygen  which  it  is  capable  of 
delivering  (cf.  p.  213). 

The  foregoing  illustrates  a  fourth  kind  of  ionic  chemical  change 
(p.  243),  namely  that  in  which  a  compound  ion  is  formed  or  decom- 
posed. Here  dichromate-ion  Cr2O7"  gives  chromic-ion  Cr""  and  water. 
For  other  illustrations  see  pp.  67,  115,  192,  195,  198,  213. 

Sulphides.  —  As  »di-basic  acid  (p.  231),  hydrogen  sulphide  gives 
both  acid  and  neutral  (or  normal)  sulphides,  such  as  NaHS  and' 
Na2S. 

The  acid  sulphides  are  obtained  by  passing  the  gas  in  excess  into 
solutions  of  soluble  bases: 

H2S  4-  JfaOH  ->  H2O  +  NaHS, 

and  are  neutral  in  reaction.  Their  negative  ion,  HS',  is  not  further 
dissociated  (see  preceding  section).  '  -^ 

By  adding  to  the  above  mentioned  solujticto.  an  amount  of  sodium 
hydroxide  equal  to  that  used  before,  and  driving,  off  the  water  by 
evaporation,  the  second  unit  of  hydrogen  is  disraaced,  and  "  neutral  " 
sodium  sulphide  is  formed: 

NaOH  +  NaHS  fc?  Na^S  +  H26  f. 

This  action  is  wholly  reversed  when  dry  sodium  sulphide  is  dissolved 
in  water,  the  salt  being  completely  hydrolyzecL  (p.  163)  to  the  acid 
salt: 

Na2St^2Na'  +  S"  1  ,  _<*;«, 

H20  ±?  OH'  +  H*  j  -^ 


254  COLLEGE    CHEMISTRY 

The  HS'  gives  a  lower  concentration  of  hydrogen-ion  than  the  water, 
and  hence  uses  up  in  its  formation  the  ions  of  hydrogen  produced  by 
the  latter,  until  an  amount  of  hydroxide-ion  equivalent  to  half  the  so- 
dium is  formed.  The  abbreviated  equation  shows  this  more  clearly  : 

S"  +  H*  +  OH'  ->  HS'  +  OH'. 

The  solution  is  therefore  strongly  alkaline  in  reaction.  In  general, 
a  "  neutral  "  salt  derived  from  an  active  base  and  a  weak  acid  is  hydro  - 
lyzed  to  some  extent  by  water  and  gives  an  alkaline  solution. 

Many  sulphides  are  insoluble  in  water,  and  these  may  be  divided 
roughly  into  three  classes: 

1.  The  sulphides  of  silver,  copper,  mercury,   and  some  other 
metals  are  exceedingly  insoluble,  and,  therefore,  do  not  interact  with 
dilute  acids  as  does  ferrous  sulphide  (p.  251).     These  may  therefore 
be  made  by  leading  hydrogen  sulphide  into  solutions  of  their  salts: 

CuS04  +  H2S  £5  CuS  J  +  H2S04. 

The  acid  produced  has  scarcely  any  effect  upon  the  sulphide,  and 
almost  no  reverse  action  is  observed.  In  this  action  the  sulphide-ion 
is  the  active  substance  and,  by  its  removal,  all  the  equilibria  (p.  251) 
are  displaced  forwards. 

2.  The  sulphides  of  iron,  zinc,  and  certain  other  metals  are  insol- 
uble in  water,  but  not  so  much  so  as  the  last  class.     Hence  they  are 
decomposed  by  dilute  acids,  and  the  reverse  of  the  above  action 
takes  place  almost  completely.     These  sulphides  must  therefore  be 
made,  either  by  combination  of  the  elements,  or  by  adding  a  soluble 
sulphide  to  a  solution  of  a  salt: 

FeSO4  +  (NH4)2S*=»FeSj  +  (NH4)2SO4. 


No  acid  is  produced  in  this  sort  of  interaction,  and  the  considerable 
insolubility  of  the  sulphide  of  iron  or  zinc  in  water  renders  the 
change  nearly  complete. 

3.  The  sulphides  of  barium,  calcium,  and  some  other  metals  (q.v.), 
although  insoluble  in  water,  are  hydrolyzed  by  it,  and  give  soluble 
products,  the  hydroxide  and  hydrosulphide  : 

2CaS  +  2H20  ±=>  Ca(OH)2  +  Ca(SH)2. 

They  may  be  prepared  by  direct  union  of  the  elements,  and  from  the 
sulphates  by  reduction  with  carbon. 


HYDROGEN    SULPHIDE  255 

The  soluble  acid  sulphides  are  oxidized  in  aqueous  solution  by 
atmospheric  oxygen: 

2NaSH  +  O2  -»  2NaOH  +  2S. 

The  sulphur  is  not  precipitated,  but  combines  with  the  excess  of  the 
sulphide,  forming  polysulphides  (see  below).  Some  sodium  thio- 
sulphate  is  produced  at  the  same  time. 

'  Polysulphides. —  When  sulphur  is  shaken  with  a  solution  of  a 
soluble  sulphide  or  acid  sulphide,  such  as  sodium  sulphide,  it  dis- 
solves, and  evaporation  of  the  solution  leaves  substances  varying  in 
composition  from  Na^  to  Na^Sg. 

When  an  acid  is  poured  into  sodium  poly  sulphide  solution,  minute 
spherules  of  rhombic  sulphur  are  precipitated: 

Na2S5  +  2HC1  ->  2NaCl  +  H2S|  +  4SJ. 

The   Chemical   Relations   of  the   Element   Sulphur.  —  In 

combination  with  metals  and  hydrogen,  sulphur  is  bivalent,  forming 
compounds  like  H2S,  FeS,  CuS,  and  HgS.  In  combination  with 
non-metals,  however,  the  valence  is  frequently  greater,  the  maximum 
being  seen  in  sulphur  trioxide,  where  the  sulphur  is  sexivalent.  Its 
oxides  are  acid-forming,  and  it  is,  therefore,  a  non-metal. 

Exercises.  —  1.  How  could  the  decomposition  of  hydrogen  sul- 
phide at  310°  be  rendered,  (a)  more  complete,  (6)  less  complete? 
Would  the  percentage  decomposed  be  affected,  (a)  by  reducing  the 
pressure,  (b)  by  mixing  the  gas  with  an  indifferent  gas? 

2.  What  are  the  relative  volumes  of  the  gases  (p.  145)  in  the  action 
of,  (a)  hydrogen  iodide  and  sulphur,  (6)  hydrogen  sulphide  and  sul- 
phur dioxide?  A 

3.  To  what  classes  of  ionic  actions  (p.  243)  do  the  interactions  of 
hydrogen  sulphide  solution  and,  (a)  oxygen  (p.  252),  (6)  sodium 
hydroxide  (p.  253),  (c)  iodine  (p.  167)  belong? 

4.  Show  which  actions  on  the  pages  referred  to  on  p.  253  illus- 
trate the  fourth  kind  of  ionic  chemical  change,  and  how  they  do  so. 

5.  Why  is  normal  sodium  sulphide  only  half  hydrolyzed  by  water? 


CHAPTER  XXII 
THE   OXIDES   AND    OXYGEN   ACIDS   OP    SULPHUR 

THE  only  important  oxides  of  sulphur  are  the  dioxide  SO,  and  the 
trioxide  S03.  They  are  the  anhydrides  (p.  51)  of  sulphurous  acid 
H;jS03  and  sulphuric  acid  H2SO4. 

The  Preparation  of  Sulphur  Dioxide  SOy  —  When  sulphur 
burns  in  air  or  oxygen,  sulphur  dioxide  is  produced  (p.  48).  The 
larger  part  of  the  sulphur  dioxide  used  in  commerce  is  probably 
obtained  by  the  roasting  (calcining)  of  sulphur  ores.  Pyrite  (FeS2), 
for  example,  burns  when  it  has  been  raised  to  the  kindling  tempera- 
ture, on  account  of  the  large  amount  of  sulphur  which  it  contains: 


4FeS3  +  HO2^2Fe2O3  +  8SO2. 

It  should  be  noted,  in  passing,  that  heating  and  roasting  or  cal- 
cining are  distinct  processes  in  chemistry.  Roasting  or  calcining 
always  assumes  the  access  of  the  air  and  employment  of  its  oxygen; 
heating,  in  the  absence  of  modifying  words,  assumes  the  exclusion  or 
the  chemical  indifference  of  the  air. 

Sulphur  dioxide  is  also  set  free  by  the  action  of  acids  upon  sul- 
phites. Sulphuric  acid  arA^i  strong  solution  of  sodium  sulphite 
may  be  used: 


+  2H?" 


H20  +SO2  (dslvd)^SO2  (gas). 


The  sulphurous  acid  being  only  moderately  ionized,  its  molecules  are 
formed  in  considerable  amount.  Being  also  very  unstable,  it  decom- 
poses spontaneously  into  water  and  sulphur  dioxide,  and  the  latter 
escapes  when  sufficient  water  for  its  solution  is  not  present. 

In  the  laboratory,  sulphur  dioxide  is  frequently  made  by  the 
reduction  of  concentrated  sulphuric  acid  by  copper  at  a  high  tem- 

256 


OXIDES    AND   OXYGEN   ACIDS   OF   SULPHUR  257 

perature.     A  part  of  the  acid  loses  oxygen  to  form  water  with  the 
hydrogen  of  another  molecule: 

H2S04->H20  +  S02(+0)  (1) 

(0)  +  H2S04  -f  Cu  -» H20  +  CuS04  (2) 

2H2S04  +  Cu  ->  2H20  +  S02  +  CuSO4 

Some  easily  oxidized  non-metals,  such  as  carbon  and  sulphur,  act  in 
the  same  way,  C  +  2H2SO4  -» 2H,O  +  2S02  +  CO2. 

Physical  and  Chemical  Properties. —  Sulphur  dioxide  is  a 
gas  possessing  a  penetrating  and  characteristic  odor.  This  is  fre- 
quently spoken  of  as  the  "odor  of  sulphur,"  but  it  should  be  remem- 
bered that  sulphur  itself  has  scarcely  any  smell  at  all.  The  weight 
of  the  G.M.V.  of  the  gas  (65.54  g.)  shows  it  to  be  more  than  twice  as 
heavy  as  air.  By  means  of  a  freezing  mixture  of  ice  and  salt,  the 
gas  is  easily  condensed  to  a  transparent  mobile  fluid,  which  boils 
at  —  8°.  The  solubility  of  the  gas  in  water  is  5000  volumes  in  100. 
The  liquid  is  completely  freed  from  the  gas  by  boiling  (cf.  p.  120). 

As  regards  chemical  properties,  sulphur  dioxide  is  stable  (p.  81). 

It  unites  with  water  to  form  sulphurous  acid,  H2S03.  Although 
the  gas  itself  sometimes  receives  this  name,  it  is  not  acid. 

Since  the  maximum  valence  of  sulphur  is  6,  sulphur  dioxide,  in 
which  but  four  of  the  valences  of  sulphur  are  used,  is  unsaturated. 
It  is  therefore  still  able  to  combine  directly  with  suitable  elements, 
such  as  chlorine  and  oxygen.  When  it  is  mixed  with  chlorine  in  sun- 
light, a  liquid,  sulphuryl  chloride  SO2C12  is  produced. 

Liquefied  sulphur  dioxide  is  now  sold  in  tin  cans,  and  is  employed 
for  bleaching  straw,  wool,  and  silk.  As  a  disinfectant  it  has  been 
displaced  to  a  large  extent  by  formaldehyde. 

Preparation  of  Sulphur  Trioxide  SO3.  —  Although  the  for- 
mation of  sulphur  trioxide  is  accompanied  by  the  liberation  of  much 
heat,  sulphur  dioxide  and  oxygen,  even  when  heated  together,  unite 
very  slowly.  Ozone,  however,  combines  with  the  former  readily. 

The  interaction  of  sulphur  dioxide  and  oxygen  is  hastened  by 
finely  divided  platinum,  which  remains  itself  unchanged  and  simply 
acts  as  a  catalytic  agent.  The  "  contact  process,"  as  this  is  called, 
has  been  rendered  availaBre  for  the  commercial  manufacture  of 
sulphur  trioxide  by  Knietsch  (1901) .  At  400°,  the  temperature  used, 
98-99  per  cent  of  the  materials  unite.  The  vaporous  product  is 


258  COLLEGE   CHEMISTRY 

condensed  by  being  led  into  97-99  per  cent  sulphuric  acid,  and  the 
concentration  of  the  liquid  is  constantly  maintained  at  this  point  by 
the  regulated  influx  of  water. 

Formerly  sulphur  trioxide  was  obtained  by  the  distillation  of 
impure  ferric  sulphate,  Fe2(SO4)3  — >  Fe2O3  +  3S03.  It  may  also  be 
prepared  by  repeated  distillation  of  concentrated  sulphuric  acid  with 
a  powerful  drying  agent,  like  phosphoric  anhydride. 

Physical  and  Chemical  Properties.  —  Sulphur  trioxide  is  a 
volatile  liquid  (b.-p.  46°).  Che  crystals,  obtained  by  cooling,  melt 
at  14.8°.  It  fumes  strongly  when  exposed  to  the  air,  in  consequence 
of  the  union  of  the  vapor  with  moisture  and  the  production  of  minute 
drops  of  sulphuric  acid.  A  white  crystalline  variety,  which  in 
appearance  closely  resembles  asbestos,  is  the  more  familiar  form  of 
the  substance.  It  has  the  formula  (S03)2. 

As  to  chemical  properties,  the  vapor  of  sulphur  trioxide  dissociates 
into  sulphur  dioxide  and  oxygen  (400°,  2%;  700°,  40%). 

Sulphur  trioxide  is  not  itself  an  acid,  but  it  is  the  anhydride  of 
sulphuric  acid.  When  placed  in  water  it  unites  vigorously,  causing 
a  hissing  noise  due  to  the  steam  produced  by  the  heat  of  the  union. 

Just  as  sulphur  trioxide  unites  with  water  to  give  hydrogen  sul- 
phate, so  it  combines  vigorously  with  many  oxides  of  metals,  producing 
the  corresponding  sulphates: 

H20  +  SO3  <=;  H,,SO4,  CaO  +  SO3  ->  CaSO4. 

The  union  of  an  oxide  of  a  non-metal  with  the  oxide  of  a  metal,  in 
this  fashion,  is  a  general  method  of  obtaining  salts  (cf.  p.  246). 

Oxygen  Acids  of  Sulphur.  —  Sulphurous  and  sulphuric  acids 
have  been  mentioned  frequ^tly  already.  Next  to  them  in  impor- 
tance come  thiosulphuric  acid  and  persulphuric  acid.  The  composi- 
tions of  the  acids  show  their  relationships: 

Hyposulphurous  acid,  H2S2O4        Thiosulphuric  acid,  H2S2O3 
Sulphurous  acid,  H2S03         Persulphuric  acid,    H2S;,O8 

Sulphuric  acid,  H2SO4 

Thiosulphuric  acid  (Gk.  0eTov,  sulphur)  is  so  named  because  it 
contains  one  unit  of  sulphur  in  place  of  ^ne  of  the  units  of  oxygen  of 
sulphuric  acid.  Besides  the  above  we  have  also  the  polythionic 
acids,  namely:  Dithionic  acid  H2S206,  trithionic  acid  H2S3Ofl,  tetra- 
thionic  acid  H2S4O6,  and  pentathionic  acid  H2S5Oe. 


OXn>F$    AND    OXYGEN   ACIDS   OF   SULPHUR  259 


SULPHURIC  ACTD  H2SO4. 

Although'*  salts  of  sulphuric  acid,  such  as  calcium  sulphate,  are 
exceedingly  plentiful  in  nature,  the  preparation  of  the  acid  by  chemi- 
cal action  upon  the  salts  is  not  practicable.  The  sulphates,  indeed, 
interact  with  all  acids,  but  the  actions  are  reversible.  The  comple- 
tion of  the  action  by  the  plan  used  in  making  hydrogen  chloride 
(p.  118),  involving  the  removal  of  the  sulphuric  acid  by  distillation, 
would  be  difficult  on  account  of  the  involatility  of  this  acid.  It  boils 
at  330°;  and  acids,  less  volatile  still,  which  might  be  used  to  liber- 
ate it,  do  not  exist.  We  are  therefore  compelled  to  build  up  sul- 
phuric acid  from  its  elements. 

The  union  of  sulphur  dioxide  and  oxygen  by  the  contact  process, 
and  combination  of  the  trioxide  with  water  (p.  257),  is  the  best 
method  for  making  a  highly  concentrated  acid.  For  obtaining 
ordinary  "  oil  of  vitriol,"  however,  the  "  chamber  process  "  is  still 
used  extensively. 

Chemistry  of  the  Chamber  Process.  —  The  gases,  the  inter- 
actions of  which  result  in  the  formation  of  sulphuric  acid,  are:  water 
vapor,  sulphur  dioxide,  nitrous  anhydride  N203  (q.v.),  and  oxygen. 
These  are  obtained,  the  first  by  injection  of  steam,  the  second  usually 
by  the  burning  of  pyrite,  the  third  from  nitric  acid,  and  the  fourth  by 
the  introduction  of.  air.  The  gases  are  thoroughly  mixed  in  large 
leaden  chambers,  and  the  sulphuric  acid  forms  droplets  which  fall 
to  the  floors.  In  spite  of  elaborate  investigations,  instigated  by  the 
extensive  scale  upon  which  the  manufacture  is  carried  on  and  the 
immense  financial  interests  involved,  some  uncertainty  still  exists  in 
regard  to  the  precise  nature  of  the  chemical  changes  which  take  place. 
According  to  Lunge  the  greater  part  of  the  product  is  formed  by  two 
successive  actions,  the  first  of  which  yields  a  complex  compound  that 
is  decomposed  by  excess  of  water  in  the  second  : 

/O-H 

H20  +  2SO,  +  N,03  4-  02  -4  2SO,  '  (1) 

"          —  JNl) 


The  group  —  NO,  nitrosyl,  is  found  in  many  compounds.     Here,  if 
it  were  displaced  by  hydrogen,  sulphuric  acid  would  result.     Hence 
this  compound  is  called  nitrosylsulphuric  add: 
O-H  OH 


260  COLLEGE    CHEMISTRY 

The  equations  (1)  and  (2)  are  not  partial  equations  for  one  inter- 
action, but  represent  distinct  actions  which  can  be  carried  out 
separately.  In  a  properly  operating  plant,  indeed,  the  nitrosyl- 
sulphuric  acid  is  not  observed.  But  when  the  supply  of  water  is 
deficient,  white  "  chamber  crystals,"  consisting  of  this  substance, 
collect  on  the  walls. 

The  explanation  of  the  success  of  this  seemingly  roundabout 
method  of  getting  sulphuric  acid  is  as  follows:  The  direct  union  of 
sulphur  dioxide  and  water  to  form  sulphurous  acid  is  rapid,  but  the 
action  of  free  oxygen  upon  the  latter,  2H2SO3  +  O2  — »  2H2S04,  is 
exceedingly  slow.  Reaching  sulphuric  acid  by  the  use  of  these  two 
changes,  although  they  constitute  a  direct  route  to  the  result,  is  not 
feasible  in  practice.  On  the  other  hand,  both  of  the  above  actions, 
(1)  and  (2),  happen  to  be  much  more  speedy,  and  so,  by  their  use, 
more  rapid  production  of  the  desired  substance  is  secured  at  the 
expense  of  a  slight  complexity. 

The  progress  of  the  first  action  is  marked  by  the  disappearance  of 
the  brown  nitrous  anhydride,  and,  on  the  introduction  of  water,  the 
completion  of  the  second  results  in  the  reproduction  of  the  same  sub- 
stance. The  nitrous  anhydride  takes  part  an  indefinite  number  of 
times  in  these  changes  and  so  facilitates  the  conversion  of  a  large 
amount  of  sulphur  dioxide,  oxygen,  and  water  into  sulphuric  acid, 
without  much  impairment  of  its  quantity. 

The  supply  of  nitrous  anhydride  is  maintained  by  the  introduction 
of  nitric  acid  vapor  into  the  chamber.  This  acid  is  made  from  con- 
centrated sulphuric  acid  and  commercial  sodium  nitrate  NaN03: 

NaN03  +  H,S04  J±  HNO3t  +  NaHSO4. 

On  account  of  the  volatility  of  the  nitric  acid,  a  moderate  heat  is 
sufficient  to  remove  it  from  admixture  with  the  other  substances,  and 
its  vapor  is  swept  along  with  the  other  gases  into  the  apparatus. 
The  initial  action  which  the  nitric  acid  undergoes: 

H2O  +  2S02  +  2HNO3  -» 2H2S04  +  N2O3, 
may  be  written,  to  show  the  anhydride  of  nitric  acid : 

H20  +  2SO2  +  H2O,N2O5  -» 2H2S04  +  N2O3. 

The  two  molecules  of  water,  one  actually,  the  other  potentially, 
present,  with  the  two  molecules  of  sulphur  dioxide,  can  furnish  two 


OXIDES   AND   OXYGEN    ACIDS   OF   SULPHUR 


261 


molecules  of  sulphurous  acid  (H2SO3).  The  N2O5  in  passing  to  the 
condition  N2O3  gives  up  the  two  units  of  oxygen  required  to  convert 
this  sulphurous  acid  into  sulphuric  acid. 

Details  of  the  Chanriber  Process.  —  The  sulphur  dioxide  is 
produced  in  a  row  of  furnaces  A  (Fig.  56).  When  good  pyrite  is 
used,  the  ore  burns  unassisted  (p.  256),  while  impure  pyrite  and 


FIG.  56. 

zinc-blende  ZnS  have  to  be  heated  artificially  to  maintain  the  com- 
bustion. The  gases  from  the  various  furnaces  pass  into  one  long 
dust-flue,  in  which  they  are  mingled  with  the  proper  proportion  of 
air,  and  deposit  oxides  of  iron  and  of  arsenic,  and  other  materials 
which  they  transport  mechanically.  From  this  flue  they  enter  the 
Glover  tower  G,  in  which  they  acquire  the  oxides  of  nitrogen.  Hav- 
ing secured  all  the  necessary  constituents,  excepting  water,  the  gases 
next  enter  the  first  of  the  lead  chambers,  large  structures  constructed 
completely  of  sheet  lead.  These  measure  as  much  as  100  X  40  X  40 
feet,  and  have  a  total  capacity  of  150,000  to  200,000  cubic  feet.  As 
the  gases  drift  through  these  chambers  they  are  thoroughly  mixed, 
and  an  amount  of  water  considerably  in  excess  of  that  actually 
required  is  injected  in  the  form  of  steam  at  various  points.  The 
acid,  along  with  the  excess  of  water,  condenses  and  collects  upon  the 
floor  of  the  chamber,  while  the  unused  gases,  chiefly  nitrous  anhydride 


262  COLLEGE    CHEMISTRY 

and  nitrogen,  the  latter  derived  from  the  air  originally  admitted, 
find  an  exit  into  the  Gay-Lussac  tower  L. 

This  is  a  tower  about  fifty  feet  in  height,  filled  with  tiles,  over 
which  concentrated  sulphuric  acid  continually  trickles.  The  object 
of  this  tower,  to  catch  the  nitrous  anhydride  and  enable  it  to  be 
reemployed  in  the  process,  is  accomplished  by  a  reversal  of  action 
(2)  above.  The  acid  which  accumulates  in  the  vessel  at  the  bottom 
of  this  tower  contains  the  nitrosylsulphuric  acid,  and  by  means  of 
compressed  air  is  forced  through  a  pipe  up  to  a  vessel  at  the  top  of 
the  Glover  tower  G.  When  this  "  nitrous  vitriol "  is  mixed  with 
dilute  sulphuric  acid  from  a  neighboring  vessel,  by  allowing  both  to 
flow  down  into  the  tower,  the  nitrous  anhydride  is  once  more  set  free 
by  the  interaction  of  the  water  in  the  dilute  acid  (action  (2)).  The 
Glover  tower  is  filled  with  broken  flint  or  tiles,  and  the  heated 
gases  from  the  furnace  acquire  in  it  their  supply  of  nitrous  anhydride. 
Their  high  temperature  causes  a  considerable  concentration  of  the 
diluted  sulphuric  acid  as  it  trickles  downward.  The  acid,  after  trav- 
ersing this  tower,  is  sufficiently  strong  to  be  used  once  more  for  the 
absorption  of  nitrous  anhydride.  To  replace  the  nitrous  anhydride 
inevitably  lost,  fresh  nitric  acid  is  furnished  by  small  open  vessels 
N,  containing  sodium  nitrate  and  sulphuric  acid,  placed  in  the  flues 
of  the  pyrite-burners.  About  4  kg.  of  the  nitrate  are  consumed  for 
every  100  kg.  of  sulphur. 

The  acid  which  accumulates  upon  the  floors  contains  but  60  to  70 
per  cent  of  sulphuric  acid,  and  has  a  specific  gravity  of  1.5-1.62. 
The  excess  of  water  is  used  to  facilitate  the  second  action.  It  is 
required  also  in  order  that  the  acid  upon  the  floor  may  not  afterwards 
absorb  and  retain  the  nitrous  anhydride,  for  this  substance  combines 
with  an  acid  containing  more  than  70  per  cent  of  hydrogen  sulphate. 

This  crude  sulphuric  acid  is  applicable  directly  in  some  chemical 
manufactures,  such  as  the  preparation  of  superphosphates  (q.v.). 
Concentration  is  effected  by  evaporation  in  pans  lined  with  lead, 
which  are  frequently  placed  over  the  pyrite-burners  in  order  to  econo- 
mize fuel.  The  evaporation  in  lead  is  carried  on  until  a  specific 
gravity  1.7,  corresponding  to  77  per  cent  concentration,  is  reached. 
Up  to  this  point  the  sulphate  of  lead  formed  by  the  action  of  the 
sulphuric  acid  produces  a  crust  which  protects  the  metal  from 
further  action.  When  a  stronger  acid  is  required,  the  water  is 
driven  out  by  heating  the  sulphuric  acid  in  vessels  of  glass  or  plati- 


OXIDES    AND   OXYGEN   ACIDS   OF   SULPHUR  263 

num,  or  even  of  cast  iron.  Commercial  sulphuric  acid,  oil  of  vitriol, 
has  a  specific  gravity  1.83-1.84,  and  contains  about  93.5  per  cent  of 
hydrogen  sulphate. 

Physical  Properties. — Pure  hydrogen  sulphate  has  a  sp.  gr.  1.85 
at  15°.  When  cooled,  it  crystallizes  (m.-p.  10°).  At  150°-180°  the 
acid  begins  to  fume,  giving  off  sulphur  trioxide.  When  boiled,  it 
yields  an  acid  of  constant  (p.  120)  boiling-point  (330°)  and  constant 
composition  (98.33  per  cent).  The  heat  of  solution  (p.  100)  of  hydro- 
gen sulphate  is  very  great  (39,170  cal.).  The  solution  is  thus  much 
more  stable  (i.e.,  it  contains  much  less  energy)  than  the  pure  sub- 
stance, and  hence  the  latter  absorbs  water  greedily.  Many  substances 
containing  hydrogen  and  oxygen  are  deprived  of  equivalent  amounts 
of  these  elements  by  sulphuric  acid.  Thus  paper,  which  is  largely 
cellulose  (C6H1005)X,  and  sugar  C12H22On  are  charred  by  it,  and 
carbon  is  set  free. 

Commercial  sulphuric  acid  is  impure.  It  contains,  for  example, 
lead  sulphate,  which  appears  as  a  precipitate  when  the  acid  is  diluted, 
as  well  as  arsenic  trioxide  and  oxides  of  nitrogen  in  combination. 

Chemical  Properties  and  Uses  of  Hydrogen  Sulphate.  —  The 

compound  is  not  exceedingly  stable,  for  dissociation  into  water 
and  sulphur  trioxide  begins  far  below  the  boiling-point,  and  is 
practically  complete  at  416°,  as  is  shown  by  the  density  of  the 
vapor.  When  raised  suddenly  to  a  red  heat  it  is  broken  up  com- 
pletely into  water,  sulphur  dioxide,  and  oxygen. 

When  sulphur  trioxide  is  dissolved  in  hydrogen  sulphate,  disul- 
phuric  acid  H2S2O7,  a  solid  compound,  is  obtained.  Hydrogen  sulphate 
containing  80  per  cent  of  disulphuric  acid  is  known  as  "  oleum,"  and 
is  employed  in  chemical  industries.  The  salts  of  disulphuric  acid 
may  be  made  by  strongly  heating  the  acid  sulphates,  for  example: 

2NaHS04  <=>  Na2S207  +  H20  f. 

In  view  of  this  mode  of  preparation  by  the  aid  of  heat,  they  are 
frequently  known  as  pyrosulphates  (Gk.  irvp,  fire).  When  they  are 
dissolved  in  water,  the  acid  sulphates  are  reproduced. 

On  account  of  the  large  quantity  of  oxygen  which  hydrogen  sul- 
phate contains,  and  its  instability  when  heated,  it  behaves  as  an 
oxidizing  agent.  This  property  has  already  been  illustrated  in  con- 
nection with  the  action  of  the  acid  upon  carbon,  sulphur,  and  copper 


264  COLLEGE   CHEMISTRY 

(p.  257),  hydrogen  iodide  (p.  251),  and  hydrogen  bromide  (p.  161). 
The  sulphuric  acid  is  in  consequence  reduced  to  sulphur  dioxide,  and 
even  to  free  sulphur  or  hydrogen  sulphide.  The  metals,  from  the 
most  active  down  to  silver  (p.  245),  are  capable  of  reducing  it,  the 
sulphates  being  formed.  Gold  and  platinum  alone  are  not  attacked. 
Free  hydrogen  itself  is  oxidized  to  water  when  passed  into  hydrogen 
sulphate  at  160°  :  S02(OH)2  +  H2  -»  S02  +  2H2O. 

With  salts  which  it  does  not- oxidize,  hydrogen  sulphate  reacts 
by  double  decomposition  and  sets  free  the  corresponding  acid.  Where 
the  new  acid  is  volatile,  as  in  the  case  of  hydrogen  chloride  (p.  117), 
we  are  furnished  with  one  of  the  cheapest  means  of  preparing  acids. 
Since  hydrogen  sulphate  is  dibasic  (p.  231),  it  forms  both  acid  and 
neutral  salts. 

Sulphuric  acid  is  used  in  almost  all  chemical  industries:  for  example, 
in  the  Le  Blanc  process  for  the  manufacture  of  soda,  in  the  refining 
of  petroleum,  in  the  manufacture  of  fertilizers,  in  the  preparation  of 
nitroglycerine  and  gun-cotton,  and  in  the  production  of  coal-tar 
dyes. 

Chemical  Properties  of  Aqueous  Hydrogen  Sulphate*  — 

The  solution  of  sulphuric  acid  H2S04,Aq  is  a  mixture,  whose  com- 
ponents are:  undissociated  molecules  H2SO4,  hydrogen-ion  H", 
hydrosulphate-ion  HSO/,  and  sulphate-ion  S04".  The  chemical 
properties  shown  by  the  solution  are  those  of  one  or  other  of  these 
components,  according  to  circumstances. 

Except  in  concentrated  solutions  (normal  or  stronger)  the  oxidiz- 
ing effects  of  the  undissociated,  molecular  substance  are  not  encoun- 
tered. 

The  presence  of  hydrogen-ion  is  shown  by  all  its  usual  properties 
(p.  232). 

Sulphate-ion  S04",  which  is  found  also  in  solutions  of  all  neutral 
and  acid  sulphates,  unites  with  all  positive  ions.  The  product,  when 
insoluble,  appears  as  a  precipitate.  The  introduction  of  barium  ions, 
for  example,  by  adding  a  solution  of  barium  nitrate  or  chloride,  is 
employed  as  a  test: 

Ba"  +  SO4"  <=>  BaS04  J,. 

Since  there  are  other  barium  salts  which  are  insoluble  in  water  (see 
Table  of  solubilities),  but  no  common  ones  which  are  not  decom- 
posed by  acids,  dilute  nitric  acid  is  first  added  to  the  solution  supposed 


OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR  265 

to  contain  the  sulphate-ion.  The  other  ions,  if  present,  then  give  no 
precipitate  with  barium-ion. 

Sulphates.  —  The  acid  sulphates,  known  also  as  bisulphates,  may 
be  produced  either  by  adding  to  sulphuric  acid  half  an  equivalent 
of  a  base,  and  evaporating:  NaOH  -f  H2SO4 t=»  H2O  +  NaHSO4,  or 
by  actions  in  which  another  acid  is  displaced,  as  in  making  hydro- 
gen chloride  (p.  117).  These  salts  are  acid  in  reaction,  as  well  as  in 
name  (cf.  p.  253), -because  HSO/, -although  a  weak,  isjaot  a  feeble 
acid.  When  heated,  they  yield  pyrosulphates  (p.  263). 

The  neutral  (or  normal)  sulphates  are  obtained  by  complete  neu- 
tralization and  evaporation,  or  by  the  second  of  the  above  methods 
when  a  sufficient  amount  of  the  salt  and  a  higher  temperature  are 

used:  NaCl  +  NaHS04  <=±  Na2S04  +  HC1| . 

They  may  also  be  made  by  precipitation,  by  oxidation  of  a  sulphide 
at  a  high  temperature,  PbS+2O2  — >  PbSO4,  or  by  addition  of  sulphur 
trioxide  to  the  oxide  of  a  metal  (p.  258). 

Normal  sulphates  of  the  heavy  metals  decompose  at  a  red  heat, 
some  giving  off  sulphur  trioxide  (p.  258),  others  sulphur  dioxide  and 
oxygen.  The  sulphates  of  the  more  active  metals  and  of  lead, 
however,  are  not  affected  by  heating. 

1 

OTHER  ACIDS  OF  SULPHUR. 

Sulphurous  Acid  H2SOs,Aq.  —  This  term  is  applied  to  the 
solution  of  sulphur  dioxide  in  water.  A  portion  of  the  sulphur 
dioxide  remains  dissolved  physically,  while  another  portion  is  in 
combination  with  the  water,  forming  sulphurous  acid.  This  in  turn 
is  ionized,  and  chiefly,  after  the  manner  of  the  weaker  dibasic  acids, 
into  two  ions,  H*  and  HS03'.  A  little  SO3"  is  formed  from  the  latter. 

Properties  of  Sulphurous  Acid.  —  The  acid  is  so  unstable  that 
it  cannot  be  obtained  excepting  in  solution  in  water.  Chemically  it 
is  a  comparatively  weak  acid.  As  a  reducing  agent,  it  is  slowly 
oxidized  by  free  oxygen,  and  rapidly  by  oxidizing  agents,  turning 
into  sulphuric  acid.  Thus,  when  free  halogens  are  added  to  the 
solution,  sulphuric  acid  and  the  hydrogen  halide  are  formed: 

H2SO3  +  H2O  +  I2  <±  H2SO4  +  2HI. 


266  COLLEGE   CHEMISTRY 

Hydrogen  peroxide,  potassium  permanganate,  and  other  oxidizing 
agents  convert  the  substance  into  sulphuric  acid  likewise. 

Sulphurous  acid  has  the  power  of  uniting  directly  with  many  organic 
coloring  matters  and,  since  the  products  of  this  union  are  usually 
colorless,  it  is  employed  as  a  bleaching  agent.  It  is  especially  useful 
with  materials  like  silk,  wool,  and  straw,  which  are  likely  to  be  de- 
stroyed by  hypochlorous  acid.  As  a  disinfectant  it  acts,  perhaps,  by 
addition  likewise.  fS 

As  a  dibasic  acid,  sulphurous  acid  forms  neutral  salts  like  Na^SOg, 
and  acid  salts  like  NaHSO3. 

Sulphites.  —  The  acid  sulphites  of  the  alkali  metals  (i.e.,  of  potas- 
sium and  sodium)  are  acid  in  reaction,  owing  to  the  appreciable  disso- 
ciation of  the  ion  HSO3'.  The  sulphites  are  readily  decomposed  by 
acids  giving  free  sulphurous  acid,  and  the  latter  partly  decomposes, 
yielding  sulphur  dioxide. 

Calcium  bisulphite  solution,  Ca(HSO3)2,  is  used  to  dissolve  the  lignin 
out  of  wood  employed  in  the  manufacture  of  paper  (q.v.). 

When  heated,  sulphites  undergo  decomposition.  The  sulphates, 
being  the  most  stable  of  all  the  salts  of  sulphur  acids,  are  formed 
when  the  salts  of  any  of  those  acids  are  decomposed  by  heating. 
The  nature  of  the  particular  salt  determines  what  other  products 
shall  appear.  Here,  one  molecule  of  the  sulphite  furnishes  three 
atoms  of  oxygen,  sufficient  to  oxidize  three  other  molecules,  and 
leaves  one  molecule  of  sodium  sulphide  behind: 

4Na2SO3  ->  Na^S  4-  3Na2S04. 


The  sulphites  are  as  readily  oxidized  as  is  the  acid  itself.  They  are 
slowly  converted,  both  in  solution  and  in  the  solid  form,  by  the  influ- 
ence of  the  oxygen  of  the  air,  into  sulphates. 

Thiosulphuric  Acid  H.2S2OZ.  —  This  acid  is  not  known  in  the  free 
condition,  but  its  salts  are  in  common  use  in  the  laboratory  and  com- 
mercially. The  sodium  salt,  for  example,  is  prepared  by  boiling  a 
solution  of  sodium  sulphite  with  free  sulphur.  The  action  is  some- 
thing like  the  addition  of  oxygen  to  sulphurous  acid: 

Na2SO3  +  S  ->  Na2S2O3    or    S03"  +  S  -»  S2O3". 
The  product,  thiosulphate  of  sodium,  is  used  in  photography  (q.v.). 


OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR  267 

By  the  addition  of  acids  to  a  solution  of  sodium  thiosulphate,  the 
thiosulphuric  acid  is  set  free,  but  the  latter  instantly  decomposes  : 

Na2S2Os  +  2HC1  <=±  H2S2OS  +  2NaCl, 


Per  sulphuric  Acid  JI2S2OS.  —  This,  like  the  other  acids  just  men- 
tioned, is  unstable,  and  can  be  kept  only  in  very  dilute  solution.  Its 
salts,  however,  are  coming  into  use  for  commercial  purposes  and  for 
"  reducing  "  negatives  in  photography.  The  salts  are  prepared  by 
electrolyzing  sodium-hydrogen  sulphate  NaHSO4  in  concentrated 
solution.  The  persulphuric  acid,  formed  by  the  union  of  the  nega- 
tive ions  in  pairs,  as  they  are  discharged  on  the  anode: 

2HSO4  +  2  0  ->  H2S208,* 

undergoes  double  decomposition  with  the  excess  of  sodium  bisulphate, 
and  the  less  soluble  sodium  persulphate  crystallizes  out.  The  other 
salts  are  made  by  double  decomposition  from  this  one. 

Compounds  of  Sulphur  and  Chlorine.  —  When  chlorine  gas  is 
passed  over  heated  sulphur  it  is  absorbed,  and  sulphur  monochloride, 
a  dark  reddish-yellow  liquid,  boiling  at  138°,  is  obtained.  The  molec- 
ular weight  of  this  substance,  as  shown  by  the  density  of  its  vapor, 
indicates  that  it  possesses  the  formula  S2C12.  When  thrown  into 
water,  it  is  rapidly  hydrolyzed,  producing  suphur  dioxide  and 
sulphur:  2S2C12  +  2H2O-*S02  +  4HC1  +  3S. 

Sulphur  itself  dissolves  very  freely  in  the  monochloride,  and  the 
solution  is  employed  in  vulcanizing  rubber. 

By  the  action  of  sulphur  dioxide  gas  upon  phosphorus  penta- 
chloride,  part  of  the  oxygen  in  the  former  is  replaced  by  chlorine: 

S02  +  PC15  -»  SOC12  +  POClj. 

The  products  are  thionyl  chloride  SOC12  and  phosphorus  oxychloride. 
The  former  is  a  colorless  liquid,  boiling  at  78°,  and  is  separated  from 
the  latter  (b.-p.  107°)  by  fractional  distillation  (see  Petroleum).  It 
is  hydrolyzed  by  water:  SOC12  +  H2O  ->  S02  +  2HC1. 

Sulphur  dioxide  and  chlorine  gases,  when  exposed  to  direct  sun- 

*  The  signs  ®  and  ©  stand  for  the  (equal)  quantities  of  electricity  carried 
by  one  equivalent  of  an  ionic  substance,  and  therefore  required  for  its  dis- 
charge and  liberation. 


268  COLLEGE,  CHEMISTRY 

light,  unite  to  form  a  liquid  known  as  sulphuryl  chloride  S02C12.  Cam- 
phor causes  the  union  to  take  place  much  more  rapidly,  owing  to  some 
catalytic  effect.  The  compound  is  a  colorless  liquid,  boiling  at  69°. 
With  water  it  gives  sulphuric  acid  and  hydrogen  chloride: 

SO2C12  +  2H20  -*  SO2(OH)2  +  2HC1. 

Exercises. —  1.  What  ground  is  there  for  assigning  the  formula 
S02  instead  of  S2O4  to  sulphur  dioxide  (p.  257)  ? 

2.  Explain  why  nitric  acid  is  completely  displaced  by  the  action  of 
sulphuric  acid  on  sodium  nitrate  (p.  260). 

3.  What  are  the  relative  volumes,  (a)  of  sulphur  dioxide  and 
nitrogen  (p.  145)  resulting  from  the  roasting  of  pyrite  (p.  256),  (6) 
of  air  and  sulphur  dioxide  in  making  sulphuric  acid,  (c)  of  nitrogen 
(left)  to  sulphur  dioxide  (used)  in  making  sulphuric  acid,  when  pyrite 
is  the  source? 

4.  Make  a  list  of,  and  classify,  the  various  applications  of  sulphuric 
acid  to  the  liberation  of  other  acids. 

5.  Formulate  the  behavior  of  the  hydrosulphate-ion  (p.  264)  when 
a  solution  of  barium  chloride  is  added  to  a  rather  concentrated  solu- 
tion of  sulphuric  acid. 

6.  Assign  to  the  proper  class  of  ionic  actions  (pp.  243,  253),  (a) 
the  action  of  iodine  on  sulphurous  acid  (p.  265),  (6)  of  sulphur  on 
sodium  sulphite  (p.  266),   (c)  the  formation  of  persulphuric  acid 
(p.  267). 


CHAPTER  XXIII 
SELENIUM    AND    TELLURIUM:    THE    PERIODIC    SYSTEM 

ALONG  with  sulphur,  chemists  group  two  other  elements,  selenium 
(Se;  at.  wt.  79.2)  and  tellurium  (Te,  at.  wt.  127.6).  If  the  nature  of 
the  chief  compounds  of  sulphur  is  kept  in  mind,  the  close  analogy 
between  the  nature  and  chemical  behavior  of  the  three  elements  and 
their  corresponding  compounds  will  be  noticed  at  once  (see  Chemical 
relations  of  the  sulphur  family,  below). 

Occurrence  and  Properties  of  Selenium,  Se.  —  Selenium 
(Gk.  o-cXrjvrj,  the  moon)  occurs  free  in  some  specimens  of  native 
sulphur,  and  in  combination  often  takes  the  place  of  a  small  part  of 
the  sulphur  in  pyrite  (FeS2).  It  is  found  free  in  the  dust-flues  of  the 
pyrite-burners  of  sulphuric  acid  works.  The  familiar  forms  are,  the 
red  precipitated  variety,  which  is  amorphous  and  soluble  in  carbon 
disulphide,  and  the  lead-gray,  semi-metallic  variety,  obtained  by 
slow  cooling  of  melted  selenium,  which  is  insoluble,  and  melts  at 
217°.  In  the  latter  form  it  has  some  capacity  for  conducting  elec- 
tricity, which  is  increased  by  exposure  to  light  in  proportion  to  the 
intensity  of  the  illumination.  It  boils  at  680°,  and  at  high  tempera- 
tures has  a  vapor  density  corresponding  to  the  formula  Se2. 

The  element  unites  directly  with  many  metals,  burns  in  oxygen  to 
form  selenium  dioxide,  and  unites  vigorously  with  chlorine. 

Compounds  of  Selenium.  —  Ferrous  selenide,  made  by  heating 
iron  filings  with  selenium,  when  treated  with  concentrated  hydro- 
chloric acid  gives  hydrogen  selenide : 

FeSe  +  2HC1  ^  H2Se  t  +  FeCl2.  ' 

The  compound  is  a  poisonous  gas,  which  possesses  an  odor  recalling 
rotten  horse-radish,  and  is  soluble  in  water.  The  solution  is  faintly 
acid  in  reaction,  and  deposits  selenium  when  exposed  to  the  action 
of  the  air  (cf.  p.  262).  Other  selenides,  which,  with  the  exception  of 

269 


270  COLLEGE    CHEMISTRY 

those  of  potassium  and  sodium,  are  insoluble  in  water,  may  be  pre- 
cipitated by  leading  the  gas  into  solutions  of  soluble  salts  of  appro- 
priate metals  (cf.  p.  254). 

The  dioxide  (SeO2)  is  a  solid  body  formed  by  burning  selenium. 
Selenious  acid  H2Se03  may  be  made  by  oxidizing  selenium  with 
boiling  nitric  acid  or  aqua  regia  (q.v.).  Unlike  sulphur  (p.  250),  the 
element  gives  little  of  the  higher  acid  H2Se04  by  this  treatment. 
The  acid  is  reduced  by  sulphurous  acid  to  selenium:  H2Se03  + 
2H2SO3  ->  2H2SO4  +  H2O  +  Se. 

No  trioxide  is  known.  Selenic  acid  H2Se04,  a  white  solid,  is  made 
by  using  the  most  powerful  oxidizing  agents  with  selenious  acid.  It 
is  itself  a  powerful  oxidizing  agent,  and,  even  in  dilute  solution, 
liberates  chlorine  from  hydrochloric  acid:  H2SeO4  +  2HC1— »H2SeO3 
-f  H2O  +  C12.  Sulphuric  acid  (cf.  p.  263),  on  the  other  hand,  is  an 
oxidizing  agent  only  in  somewhat  concentrated  form,  and  even  then 
it  can  oxidize  hydrobromic  acid  (p.  161),  but  not  hydrochloric  acid. 

Tellurium  Te.  —  Tellurium  (Lat.  tellus,  the  earth)  occurs  in 
sylvanite  in  combination  with  gold  and  silver.  It  is  a  white,  metallic, 
crystalline  substance,  melting  at  452°.  When  formed  by  precipita- 
tion it  is  a  black  powder.  The  free  element  unites  with  metals 
directly,  and  burns  in  air  to  form  the  dioxide. 

The  compounds  of  tellurium  are  similar  in  composition  and  mode 
of  preparation  to  those  of  selenium.  Some  differences  in  chemical 
behavior  are  significant,  however.  Thus,  tellurious  acid  H2TeO3  is 
a  very  feeble  acid  and  is  also  somewhat  basic,  a  sulphate  (2Te02,SO3) 
and  a  nitrate  (Te2O3(OH)NO3)  being  known.  In  this  respect  it 
differs  markedly  from  sulphurous  acid.  Telluric  acid  does  not  affect 
indicators,  and  is  therefore  actually  more  feebly  acidic  than  is 
hydrogen  sulphide.  Tellurium  tetrachlorideTeC!4,  although  hydrolyzed 
by  water,  exists  in  solution  with  excess  of  hydrogen  chloride: 
TeCl4  +  3H2O  <=±  H2Te03  +  4HC1,  showing  the  tellurious  acid  to  be 
basic  in  properties  and  the  element  tellurium  to  be,  to  a  certain 
degree,  a  metal. 

The    Chemical    Relations    of    the    SulpJiur    Family.  —  It 

will  be  seen  that  sulphur,  selenium,  and  tellurium  are  bivalent  ele- 
ments when  combined  with  hydrogen  or  metals.  In  combination 
with  oxygen  they  form  unsaturated  compounds  of  the  form  XIV  O2, 


TITK    PERIODIC    SYSTEM  271 

while  their  highest  valence  is  found  in  SO3,  TeO3,  and  H2SeO4,  where 
they  must  be  sexivalent.  The  general  behavior  of  corresponding 
compounds  is  very  similar.  At  the  same  time,  there  is  in  all  cases 
a  progressive  change  as  we  proceed  from  sulphur  through  selenium 
to  tellurium.  The  elementary  substances  themselves,  for  example, 
become  more  like  metals,  physically,  and  they  show  higher  and 
higher  melting-points.  The  affinity  for  hydrogen  decreases,  as  is 
shown  by  the  increasing  ease  with  which  the  compounds  H2X  are 
oxidized  in  air.  The  affinity  for  oxygen  likewise  decreases,  for  the 
elements  become  increasingly  difficult  to  raise  to  the  highest  state  of 
oxidation.  On  the  other  hand,  the  tendency  to  form  higher  chlo- 
rides becomes  greater.  We  note  also  that  the  compounds  H2XO4 
become  less  and  less  active  as  acids,  and  that  a  basic  tendency  begins 
to  assert  itself. 

THE  PERIODIC  SYSTEM. 

Classification,  or  the  arrangement  of  facts  on  the  basis  of  likeness, 
is  part  of  the  method  of  science.  It  is  needed  to  make  possible  the 
systematic  description  of  the  ascertained  facts,  and  to  furnish  a 
guide  in  investigation,  by  suggesting  relations  and  so  pointing  out 
directions  in  which  new  facts  of  interest  may  be  found.  Thus,  we 
have  treated  the  halogens  as  a  group;  and  chemists,  knowing  how 
hypochlorous  acid  (HC10)  and  perchloric  acid  (HC1O4),  and  their 
salts,  are  made,  have  been  led  to  attempt  to  obtain  related  substances, 
like  HIO  and  HBr04  and  their  salts,  by  methods  suggested  by 
analogy. 

Metallic  and  Non-Metallic  Elements.  —  Thus  far  we  have 
found  the  division  into  metallic  and  non-metallic  elements  very  ser- 
viceable for  classification  in  terms  of  chemical  relations  (p.  116). 
This  distinction  we  shall  continue  to  employ.  The  metallic,  or  posi- 
tive elements  (p.  82),  (1)  form  positive  radicals  and  ions  containing 
no  other  element  (cf.  p.  233).  Thus  the  metals  give  sulphates, 
nitrates,  carbonates,  and  other  salts,  which  furnish  a  metallic  ion 
together  with  the  ions  SO4",  NO/,  and  C03".  (2)  Their  hydroxides, 
KOH,  Ca(OH)2,  etc.,  give  the  same  metallic  ion,  and  the  rest  of  the 
molecule  forms  hydroxide-ion.  That  is  to  say,  their  hydroxides  are 
bases  and  their  oxides  are  basic.  The  metallic  elements  often  enter 
with  other  elements  into  the  composition  of  a  negative  ion,  as  is  the 


272  COLLEGE    CHEMISTRY 

case  with  manganese  in  K.MnO4,  with  chromium  in  K2.Cr2O7,  and 
with  silver  in  K.Ag(CN)2. 

The  non-metallic  or  negative  elements,  (1)  are  found  chiefly  in 
negative  radicals  and  ions.  They  form  no  nitrates,  sulphates,  car- 
bonates, etc.,  for  they  could  not  do  so  without  themselves  alone 
constituting  the  positive  ion.  We  have  no  such  salts  of  sulphur, 
carbon,  or  phosphorus,  for  example.  (2)  Their  hydroxides,  like 
C102OH,  P(OH)3,  SO2(OH)2,  furnish  no  hydroxyl  ions,  as  this  would 
involve  the  same  consequence.  These  hydroxides  are  divided  by 
dissociation,  in  fact,  so  that  the  non-metal  forms  part  of  a  compound 
negative  radical,  and  the  other  ion  is  hydrogen-ion,  C1O3.H,  PO3H.H2, 
S04.H2.  Their  oxides  are  acidic.  (3)  Their  halogen  compounds, 
like  PBr3  (p.  163)  and  S2C12  (p.  267),  are  completely  hydrolyzed  by 
water,  and  the  actions  are  not,  in  general,  reversible.  The  halides 
of  the  typical  metals  are  not  hydrolyzed  (see  Chap,  xxxii) . 

The  "distinction  is  not  perfectly  sharp,  however.  Thus,  zinc  (q.v.) 
gives,  both  salts  like  the  sulphate,  Zn.S04,  and  chloride,  Zn.Cl2,  and 
compounds  like  sodium  zincate  (p.  67),  ZnO^Na^. 

Classification  by  Atomic  Weights.  —  Newlands  (1863-4)  dis- 
covered a  surprising  regularity  that  became  apparent  when  the  ele- 
ments were  placed  in  the  order  of  ascending  atomic  weight.  Omitting 
hydrogen  (at.  wt.  1)  the  first  seven  were:  lithium  (7),  glucinum  (9), 
boron  (11),  carbon  (12),  nitrogen  (14),  oxygen  (16),  fluorine  (19). 
These  are  all  of  totally  different  classes,  and  include  first  a  metal 
forming  a  strongly  basic  hydroxide,  then  a  metal  of  the  less  active 
sort,  then  five  non-metals  of  increasingly  negative  character,  the  last 
being  the  most  active  non-metal  known.  The  next  element  after 
fluorine  (19)  was  sodium  (23),  which  brings  us  back  sharply  to  the 
elements  that  form  strongly  basic  hydroxides.  Omitting  none,  the 
next  seven  elements  were:  sodium  (23),  magnesium  (24.4),  aluminium 
(27),  silicon  (28.4),  phosphorus  (31),  sulphur  (32),  chlorine  (35.5). 
In  this  series  there  are  three  metals  of  diminishing  positiveness, 
followed  by  four  non-metals  of  increasing  negative  activity,  the  last 
being  a  halogen  very  like  fluorine.  On  account  of  the  fact  that  each 
element  resembles  most  closely  the  eighth  element  beyond  or  before 
it  in  the  list,  the  relation  was  called  the  law  of  octaves.  After  chlorine 
the  octaves  become  less  easy  to  trace. 

That  this  periodicity  in  chemical  nature  is  more  than  a  coincidence 


T1IK    PKRlODir    S  VST  KM  273 

is  shown  by  the  fact  that  the  valence  and  even  the  physical  proper- 
ties, such  as  the  specific  gravity,  show  a  similar  fluctuation  in  each 
series.  In  the  first  two  series  the  compounds  with  other  elements 
are  of  the  types: 

LiCl,  Old,,  BC13(  CCU      {S£oH,,FH. 
NaCl,  MgCL,  A1C13,  SiCl4, 

Thus  the  valence  towards  chlorine  or  hydrogen  ascends  to  four  and 
then  reverts  to  one  in  each  octave.     The  highest  valence,  shown  in 
oxygen  compounds,  ascends  from  lithium  to  nitrogen  with  values  one 
to  five,  and  then  fails  because  compounds  are  lacking.     In  the 
second  octave,  however,  it  goes  up  continuously  from  one  to  seven. 
Again,  the  specific  gravities  of  the  elements  in  the  second  series, 
using  the  data  for  red  phosphorus  and  liquid  chlorine,  are: 
Na  0.97,  Mg  1.75,  Al  2.67,  Si  2.49,  P  2.14,  S  2.06,  Cl  1.33. 

Mendelejeff's  Scheme In  1869  Mendelejeff  published  an 

important  contribution  towards  adjusting  the  difficulty  which  the  ele- 
ments following  chlorine  presented,  and  developed  the  whole  concep- 
tion so  completely  that  the  resulting  system  of  classification  has  been 
connected  with  his  name  ever  since.  Almost  simultaneously  Lothar 
Meyer  made  similar  suggestions,  but  did  not  urge  them  with  the 
same  conviction  or  elaborate  them  so  fully.  The  table  on  the  fol- 
lowing page,  in  which  the  atomic  weights  are  expressed  in  round 
numbers,  is  a  modification  of  one  of  Mendelejeff  s. 

The  chief  change  from  the  arrangement  in  simple  octaves  is  that 
the  third  series,  beginning  with  potassium,  is  made  to  furnish  mate- 
rial for  two  octaves,  potassium  to  manganese  and  copper  to  bromine, 
and  is  called  a  long  series.  The  valences  fall  in  with  this  plan  fairly 
well.  Copper,  while  usually  bivalent,  forms  also  a  series  of  com- 
pounds in  which  it  is  univalent.  Iron,  cobalt,  and  nickel  cannot  be 
accommodated  in  either  octave,  as  their  valences  are  always  two  or 
three.  At  the  time  Mendelejeff  made  the  table,  three  places  in  the 
third  series  had  to  be  left  blank,  as  a  trivalent  element  [Sc]  was  lack- 
ing in  the  first  octave,  and  a  trivalent  [Ga]  and  a  quadrivalent  one 
[Ge]  in  the  second.  These  places  have  since  been  filled,  as  we  shall 
presently  see.  The  first  two,  (the  short)  series  have  been  split  in  the 


274 


COLLEGE    CHEMISTRY 


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TMK    PERK  JDK1    SYSTEM  275 

table,  as  lithium  and  sodium  closely  resemble  potassium,  while  the 
remaining  members  of  these  series  fall  more  naturally  over  the  corre- 
sponding elements  of  the  second  octave  of  the  third  series. 

The  fourth  series  (long)  is  nearly  complete.  It  begins  with  an 
active  alkali  metal,  rubidium,  and  ends  with  iodine,  a  halogen.  The 
rule  of  valence  is  strictly  preserved  throughout  the  series,  and  in  gen- 
eral the  elements  fall  below  those  which  they  most  closely  resemble. 

The  fifth,  sixth,  and  seventh  (long)  series  are  incomplete,  but  the 
order  of  the  atomic  weights  and  the  valence  enable  us  satisfactorily 
to  place  those  elements  which  are  known.  The  chemical  relations  to 
elements  of  the  fourth  series  justify  the  position  assigned  to  each. 
Caesium,  for  example,  is  the  most  active  of  the  alkali  metals; 
barium  has  always  been  classed  with  strontium,  and  bismuth  with 
antimony. 

In  two  cases  a  slight  displacement  of  the  order  according  to  atomic 
weights  is  necessary.  Cobalt  is  put  before  nickel  because  it  resembles 
iron  more  closely.  Tellurium  and  iodine  are  placed  in  that  order  to 
bring  them  into  the  sulphur  and  halogen  groups  respectively.  Their 
valence  and  other  chemical  relations  both  require  this.  The  general 
agreement,  however,  is  very  remarkable. 

General  Relations  in  the  System.  —  In  every  octave  the  val- 
ence towards  oxygen  ascends  from  one  to  seven,  while  that  towards 
hydrogen,  in  the  cases  of  the  last  four  elements  (when  they  combine 
with  hydrogen  at  all),  descends  from  four  to  one.  The  physical 
properties  fluctuate  within  the  limits  of  each  series  in  a  similar  way. 
The  values  of  each  physical  constant  for  corresponding  members  of 
the  successive  series  do  not  exactly  coincide,  however.  A  progres- 
sive change,  as  we  descend  each  vertical  column,  is  the  rule.  Thus 
the  specific  gravities  (water  =  1)  of  the  alkali  metals  rise  from 
lithium  (0.53)  to  caesium  (1.87).  In  the  same  group  the  melting- 
points  descend  from  lithium  (186°)  to  caesium  (26.5°). 

As  yet  no  exact  mathematical  relation  between  the  values  for  any 
property  and  the  values  of  the  atomic  weights  has  been  discovered; 
only  a  general  relationship  can  be  traced.  Anticipating  the  discov- 
ery of  some  more  exact  mode  of  stating  the  relationship  in  each  case, 
and  remembering  that  similar  values  of  each  property  recur  periodi- 
cally, usually  at  intervals  corresponding  to  the  length  of  an  octave 
or  series,  the  principle  which  is  assumed  to  underlie  the  whole,  the 


2TO  COLLEGE    CHEMISTRY 

periodic  law,  is  stated  thus:  All  the  properties  of  the  elements  are 
periodic  functions  of  their  atomic  weights. 

That  the  chemical  relations  of  the  elements  vary  just  as  do  the 
physical  properties  of  the  simple  substances  is  easily  shown.  Thus, 
each  series  begins  with  an  active  metallic  (positive)  element,  and 
ends  with  an  active  non-metallic  (negative)  element,  the  intervening 
elements  showing  a  more  or  less  continuous  variation  between  these 
limits.  Again,  the  elements  at  the  top  are  the  least  metallic  of  their 
respective  columns.  As  we  descend,  the  members  of  each  group  are 
more  markedly  metallic  (in  the  first  columns),  or,  what  is  the  same 
thing,  less  markedly  non-metallic  (in  the  later  columns;  cf.  p.  271). 

In  the  first  series  boron  is  the  first  non-metal  we  encounter.  In 
the  second  series  silicon  is  the  first  such  element.  In  the  third  there 
is  more  difficulty  in  deciding.  Titanium,  vanadium,  and  germanium 
are  usually,  though  with  questionable  propriety,  classed  as  metals. 
Selenium  is  undoubtedly  a  non-metal.  Arsenic  is,  on  the  whole,  a 
non-metal.  In  the  fourth  series  tellurium  is  commonly  considered 
to  be  the  first  non-metal.  Thus  a  zigzag  line,  shown  in  the  table, 
separates  all  the  non-metals  from  the  rest  of  the  elements,  and  con- 
fines them  in  the  right-hand  upper  corner. 

A  more  compact  form  of  the  table  is  printed  at  the  end  of  this 
book,  opposite  the  rear  board.  The  only  difference  between  this 
and  the  other  is  that  the  two  octaves  of  each  long  series  have  been 
placed  in  the  same  set  of  seven  main  columns.  The  iron,  palladium, 
and  platinum  groups  occupy  a  column  on  the  right  of  the  main 
columns,  and  are  often  called  collectively  the  eighth  group.  The 
newly  discovered  elements,  found  chiefly  in  the  air,  have  been  placed 
at  the  left-hand  side.  Since  they  do  not  enter  into  combination  at 
all,  their  valence  may  appropriately  be  given  as  zero.  With  the 
exception  of  argon,  the  values  of  their  atomic  weights  agree  well  with 
this  assignment.  Hydrogen  is  the  only  common  element  whose 
place  is  still  in  debate.  The  valence  is  shown  by  the  general  formulae 
at  the  head  of  each  column. 

Applications  of  the  Periodic  System.  — The  system  has  found 
application  chiefly  in  four  ways: 

1.  In  the  prediction  of  new  elements.  Mendelejeff  (1871)  drew 
attention  to  the  blank  then  existing  between  calcium  (40)  and  tita- 
nium (48).  He  predicted  that  an  element  to  fit  this  place  would 


THE   PERIODIC   SYSTEM  277 

have  an  atomic  weight  44  and  would  be  trivalent.  From  the  nature 
of  the  surrounding  elements,  he  very  cleverly  deduced  many  of  the 
physical  and  chemical  properties  of  the  unknown  element  and  of  its 
compounds.  In  1879  Nilson  discovered  scandium  (44),  and  its 
behavior  corresponded  closely  with  that  predicted.  Mendelejeff 
described  accurately  two  other  elements,  likewise  unknown  at  the 
time.  In  1875  Lecoque  de  Boisbaudran  found  gallium,  and  in  1888 
Winkler  discovered  germanium,  and  these  blanks  were  filled. 

2.  By  enabling  us  to  decide  on  the  correct  values  for  the  atomic 
weights  of  some  elements,  when  the  equivalent  weights  have  been 
measured,  but  no  volatile  compound  is  known  (cf.  pp.  130  and  147). 
Thus,  the  equivalent  weight  of  indium  was  38  and,  as  the  element 
was  supposed  to  be  bivalent,  it  received  the  atomic  weight  76.     It 
was  quite  out  of  place  near  arsenic  (75),  however,  being  decidedly  a 
metal.     As  a  trivalent  element  with  the  atomic  weight  115,  it  fell 
between  cadmium  and  tin.     Later  work  fully  justified  the  change. 
Quite  recently,  radium  (q.v.)  has  been  discovered,  and  found  to  have 
the  equivalent  weight  113.25  and  to  resemble  barium.     If,  like 
barium,  it  is  bivalent,  it  occupies  a  place  under  this  element,  in  the 
last  series. 

3.  By  suggesting  problems  for  investigation.     The  periodic  system 
has  been  of  constant  service  in  the  course  of  inorganic  research,  and 
has  often  furnished  the  original  stimulus  to  such  work  as  well.     For 
example,  the  atomic  weight  of  tellurium  bore  the  value  128  when  the 
table  was  first  constructed,  and  it  was  confidently  expected  that 
reexamination  would  bring  this  value  below  that  of  iodine  (then  127, 
now  126.97).     Several  most  careful  studies  of  the  subject  have  been 
made  by  different  methods.     It  seems  probable  that  the  real  value 
of  the  atomic  weight  is  not  far  from  Te  =  127.6,  and  therefore  more 
than  half  a  unit  greater  than  that  of  iodine.     Since,  however,  mathe- 
matical correspondence  is  found  nowhere  in  the  system,  the  existence 
of  marked  inconsistencies  like  this  need  not  shake  our  confidence  in 
its  value  when  it  is  used  with  due  consideration  of  the  degree  of 
correspondence  to  be  expected. 

4.  By  furnishing  a  comprehensive  classification  of  the  elements, 
arranging  them  so  as  to  exhibit  the  relationships  among  the  physical 
and  chemical  properties  of  the  elements  themselves  and  of  their  com- 
pounds.   Constant  use  will  be  made  of  this  property  of  the  table  in 
the  succeeding  chapters.     Having  disposed  of  the  halogen  and  sul- 


278  COLLEGE    CHEMISTRY 

phur  families,  situated,  respectively,  in  the  seventh  and  sixth  columns 
of  the  table  (at  the  end  of  this  book),  we  shall  next  take  up  nitrogen 
and  phosphorus  from  the  right  side  of  the  fifth  column.  Then  from 
the  fourth  column,  we  shall  select  carbon  and  silicon,  and  from  the 
third  boron,  leaving  the  other,  more  decidedly  metallic  elements  for 
later  treatment. 

Exercises.  —  1.   Can  you  explain  the  presence  of  free  selenium  in 
the  flues  of  pyrite-burners  (p.  270)  ? 

2.  How  should  you  attempt  to  obtain  KIO  and  KBrO4? 

3.  Make  a  list  of  bivalent  elements  and  criticize  this  method  of 
grouping  as  a  means  of  chemical  classification. 

4.  Write  down  the  symbols  of  the  elements  in  the  fourth  series 
(that  beginning  with  rubidium,  and  ending  with  iodine)  on  p.  274. 
Record  the  valence  of  each  element  toward  oxygen,  using  for  refer- 
ence the  chapters  in  which  the  oxygen  compounds  are  described. 


CHAPTER  XXIV 
NITROGEN    AND    AMMONIA 

WHEN  the  oxygen  of  the  air  was  removed  (p.  45),  a  gas  remained 
which  was  largely  nitrogen.  It  did  not  support  combustion  or  life, 
and  was  named  azote  (Gk.  £omKos,  life).  The  English  name  records 
the  fact  that  it  is  an  important  constituent  of  saltpeter  KNO3  (Lat. 
nitrum). 

Tlie    Chemical   Relations   of   the   Element   Nitrogen.  —  In 

compounds  with  hydrogen  and  the  metals  nitrogen  is  trivalent, 
while  in  those  containing  oxygen  and  other  negative  elements,  it  is 
frequently  quinquivalent.  It  is  a  non-metal,  for  its  oxides  are  acidic 
(p.  272).  The  compounds  of  nitrogen  are  often  extremely  active 
and  interesting.  Those  of  them  which  we  have  to  discuss  in  inor- 
ganic chemistry  are  ammonia  NH3  and  nitric  acid  HN03,  and 
several  related  substances. 

Occurrence  and  Preparation.  —  Free  nitrogen  is  present  ir 
the  air.  The  nitrates  of  potassium  and  sodium  are  found  in  Bengal 
and  Peru  respectively.  Natural  manures,  such  as  guano,  contain 
large  quantities  of  nitrogen  compounds,  and  owe  part  of  their  value 
as  fertilizers  to  this  fact.  Nitrogen  is  an  essential  constituent  of 
vegetable  and  animal  matter.  The  albumins,  for  example,  contain 
on  an  average  about  15  per  cent  of  combined  nitrogen. 

Nitrogen  containing  about  one  per  cent  of  argon  (q.v.)  is  obtained 
by  burning  phosphorus  in  air,  or  by  passing  air  over  heated  copper. 

Pure  nitrogen  is  prepared  by  heating  ammonium  nitrite : 

NH4NO2  ->  2H20  +  N2. 

In  practice,  strong  solutions  of  ammonium  chloride  and  sodium 
nitrite  are  mixed,  a  double  decomposition  results  in  the  formation  of 
ammonium  nitrite,  NH4C1  +  NaN02  <=»  NH4NO2  +  NaCl,  and  this 
breaks  up  when  heat  is  applied,  giving  nitrogen. 

279 


280  COLLEGE    CHEMISTRY 

We  may  also  prepare  nitrogen  by  the  oxidation  of  ammonia  (NH3) , 
by  passing  the  latter  over  heated  cupric  oxide  (see  p.  282),  or  by  the 
reduction  of  nitric  oxide  (NO)  by  passing  this  gas  over  heated 
copper. 

Physical  and  Chemical  Properties.  —  Nitrogen  is  a  colorless, 
tasteless,  odorless  gas,  as  we  should  expect  from  the  fact  that  air 
possesses  these  properties.  It  forms  a  colorless  liquid,  boiling  at 
-  194°,  and  a  white  solid  (m.-p.  -  214°).  The  solubility  in  water 
(1.6  vols.  in  100)  is  less  than  that  of  oxygen. 

The  density  of  the  gas  shows  the  formula  of  free  nitrogen  to 
be  N2. 

Nitrogen  unites  with  few  elements  directly.  At  ordinary  tempera- 
tures it  is  almost  absolutely  indifferent.  When  passed  over  heated 
lithium,  calcium,  magnesium,  or  boron,  it  forms  nitrides,  in  which  it 
is  trivalent.  These  have  the  formulae  Li3N,  Ca3N2,  Mg3N2,  and  BN 
respectively.  When  the  gas  is  mixed  with  oxygen  or  hydrogen,  and 
sparks  from  an  induction  coil  are  passed  between  platinum  wires 
through  the  mixtures,  small  amounts  of  nitrogen  tetroxide  N2O4, 
and  ammonia  NH3,  respectively,  are  produced. 

One  case  of  direct  union  of  nitrogen  is  of  economic  importance. 
The  supply  required  by  most  plants  is  obtained  from  nitrogen  com- 
pounds contained  in  fertilizers,  or  equivalent  substances  already 
present  in  the  soil.  With  the  leguminosce  (peas,  beans,  clover,  etc.), 
however,  are  found  associated  certain  bacteria,  which  nourish  in 
nodules  upon  their  roots.  These  bacteria  have  the  power  of  taking 
free  nitrogen  from  the  air,  which  penetrates  the  soil,  and  producing 
compounds  containing  nitrogen.  The  nodules  often  contain  over 
five  per  cent  of  combined  nitrogen.  The  compounds  are  chiefly 
albumins,  which  are  afterwards  digested  and  absorbed  by  the  roots 
of  the  plant. 

Compounds  of  Nitrogen  and  Hydrogen. — The  commonest  and 
longest  known  of  these  substances  is  ammonia  NH3,  which  was  first 
described  by  Priestley  (1774)  and  named  "  alkaline  air."  Curtius 
discovered  hydrazine  N2H4  in  1889,  and  hydrazoic  acid  HN3  in  1890. 
Hydroxylamine  NH3O,  discovered  by  Lessen  in  1865,  is  similar  to 
ammonia  in  chemical  behavior. 


NITROGEN   AND   AMMONIA  281 

AMMONIA  NH3. 

Preparation.  —  1.  When  sparks  from  an  induction  coil  are  passed 
through  a  mixture  of  nitrogen  and  hydrogen,  very  small  proportions 
of  the  materials  unite  to  form  ammonia  (see  below). 

2.  When  water  is  added  to  the  nitride  of  magnesium  or  calcium 
(p.  280),  ammonia  is  given  off  and  the  hydroxide  of  the  metal 

Mg3N2  -f  6H2O  ->  3Mg(OH)2  +  2NH3. 

3.  When  parts  of  animals,  particularly  the  horns,  hides,   and 
feathers,  which  contain  complex  compounds  of  carbon,  nitrogen, 
hydrogen,  and  oxygen,  are  heated  strongly,  much  of  the  nitrogen  is 
driven  out  as  ammonia.     Hence  the  name  spirit  of  hartshorn,  applied 
to  the  aqueous  solution.  a 

4.  The  entire  commercial  supply  is  obtained  as  a  by-product  from 
operations  like  the  manufacture  of  illuminating-gas  and  of  coke,  in 
which  the  destructive  distillation  of  coal  takes  place.     The  crude 
mixture  of  gases  passes  first  through  water,  in  which  most  of  the 
ammonia  dissolves. 

5.  In  the  laboratory,  a  mixture  of  slaked  lime  and  some  salt  of 
ammonium,  such  as  ammonium  chloride,  either  with  or  without 
water,  is  heated  in  a  flask  or  retort  provided  with  a  delivery  tube: 

Ca(OH)2  +  2NH4C1  ±?  CaCl2  +  2NH4OH, 
NH4OH  t+  H2O    +  NH3. 

6.  Warming  the  aqueous  solution  gives  a  steady  stream  of  the  gas. 
Since  the  gas  is  very  soluble  in  water,  it  is  collected  over  mercury 

or  by  upward  displacement  of  air.     It  is  dried  with  quicklime. 

Physical  Properties.  —  Ammonia  is  a  colorless  gas  with  a  pun- 
gent, characteristic  odor  familiar  in  smelling-salts.  The  G.M.V.  of 
the  gas  weighs  17.26  g.,  so  that  the  density  is  little  more  than  half 
that  of  air  (cf.  p.  126).  When  liquefied  it  boils  at  -34°  and  the  solid 
is  white  and  crystalline  (m.-p.  —  77°).  One  volume  of  water  dis- 
solves 1148  volumes  of  the  gas  at  0°,  764  volumes  at  16°,  and  306 
volumes  at  50°.  The  35  per  cent  solution,  sold  as  "  concentrated 
ammonia/'  has  a  sp.  gr.  0.882.  The  whole  of  the  dissolved  gas  may 
be  removed  by  boiling  (cf.  p.  120). 

Liquefied  ammonia  is  used  in  refrigeration.  The  gas  is  liquefied 
by  compression,  and  the  heat  which  is  thus  liberated  is  removed  by 


282  COLLEGE   CHEMISTRY 

flowing  water  which  surrounds  the  pipes.  The  liquid  then  passes  into 
other  pipes  immersed  in  calcium  chloride  brine,  and  is  there  allowed 
to  evaporate,  the  gas  returning  to  the  compressor.  The  heat  of 
vaporization,  260  calories  per  gram,  is  taken  from  the  brine,  which  is 
thus  partially  frozen.  The  resulting  freezing  mixture  of  ice  and 
calcium  chloride  solution  is  then  distributed  to  the  localities  to  be 
cooled.  The  ammonia  and  brine  remain  within  their  respective 
closed  systems  of  pipes,  and  are  used  over  and  over  again. 

Chemical  Properties.  —  The  discharge  from  an  induction-coil 
decomposes  ammonia  (to  the  extent  of  94-98  per  cent)  into  nitrogen 
and  hydrogen.  Under  the  same  circumstances,  union  of  the  con- 
stituents also  occurs  (up  to  2-6  per  cent) : 

2NH3  fc?  N2  +  3H2. 

Ammonia  reduces  certain  oxides,  when  led  over  them: 
3CuO  +  2NH3  ->  3Cu  +  3H2O  +N2, 

and  burns  in  pure  oxygen  with  the  same  result. 

Chlorine  and  bromine  combine  with  the  hydrogen  and  liberate  the 
nitrogen  of  ammonia.  This  action  may  be  used  for  obtaining  a 
stream  of  nitrogen,  provided  excess  of  chlorine  is  avoided  (see  Nitro- 
gen trichloride,  below).  Chlorine  is  led  into  a  solution  of  ammonium 

2NH4C1  +  3C12  -»  N2 1  +  8HC1. 

The  power  to  form  a  base  by  combination  with  water  is  the  most 
characteristic  property  of  ammonia: 

NH3  (gas)  £=?  NH3  (dslvd)  +  H20  <=±  NH4OH  t?  NH4'  +  OH'. 

Probably  only  a  small  proportion  of  the  gas  is  actually  combined  at 
any  one  time,  the  greater  part  being  simply  dissolved. 

The  gas  unites  also  with  acids,  forming  salts  (cf.  p.  232),  which,  in 
solution,  are  highly  ionized: 

NH3  4-  HC1      ->  NH4.C1, 
NH3  +  HN03  ->  NH4.NO3. 

Ammonium  Compounds.  —  Since  NH4  plays  the  part  of  a 
metallic  element,  entering  into  the  composition  of  a  base  and  of  a 
series  of  salts,  it  is  named  ammonium.  As  this  radical  forms  a  univa- 


NITROGEN   AND    AMMONIA  283 

lent,  positive  ion  and  gives  a  distinctly  alkaline  base,  it  is  classed  with 
the  metallic  elements  of  the  alkalies  (q.v.). 

Ammonium  hydroxide,  although  less  completely  ionized  than  potas- 
sium hydroxide,  affects  litmus  easily.  In  a  normal  solution  about 
0.4  per  cent  of  the  ammonia  is  in  the  form  of  ammonium-ion  NH4". 
When  an  acid  is  added  to  the  solution,  the  correspondingly  small 
amount  of  hydroxide-ion  which  exists  in  it  is  removed  and  the 
various  equilibria  are  displaced  forwards.  The  final  result  is  the 
same  as  with  any  other  base: 

NH3(dslvd)  +  H20  <=>  NH4OH  <=?  NH4'  +  OH'  K 

HCII=>   ci  +  H-  y-  H'a 

When  strongly  heated,  all  ammonium  salts  are  decomposed  and, 
usually,  give  ammonia  and  the  acid.  When  the  latter  is  volatile, 
the  whole  material  of  the  salt  is  thus  converted,  into  gas.  If  the 
acid  is  volatile  without  permanent  decomposition,  it  reunites  with 
the  ammonia  when  the  vapor  is  cooled: 

NH4C1  <=»  HC1  +  NH3. 

The  use  of  ammonium  chloride  (salammoniac)  in  soldering  depends 
on  the  dissociation  of  the  salt,  by  the  heat  of  the  iron,  and  the  action 
of  the  liberated  hydrochloric  acid  on  the  oxide  which  covers  the 
surface  of  the  metal  to  be  soldered. 

The  test  for  ammonium  salts  is  to  warm  them,  dry  or  in  solution, 
with  a  base: 

(NH4)2S04  ±=;  S04"  +  2NH4*  )  _  2NH  OH  -»  2H  O 
2KOH  s=?  2K*    +  2OH'  J  -"  ^M4U±1  *=?  zti^ 

when  the  odor  of  ammonia  becomes  noticeable.  When  the  solution 
is  used,  it  is  the  tendency  of  the  NH4*  and  OH'  to  unite  to  form  the 
slightly  ionized  molecular  hydroxide  that  sets  the  other  equilibria 
in  motion. 

In  ammonium  salts,  the  nitrogen  is  quinquivalent. 


Hydrazine  N^H^.  —  By  reduction  of  a  compound  of  nitric  oxide 
and  potassium  sulphite  by  means  of  sodium  amalgam,*  a  solution 
of  hydrazine  hydrate  is  obtained: 

K2SO3,2NO  +  3H2  -»  N2H4,H2O  +  K2SO4. 

*  The  sodium  dissolved  in  the  mercury  interacts  with  the  water,  giving 
hydrogen  (see  Active  state  of  hydrogen)  . 


284  COLLEGE   CHEMISTRY 

When  the  hydrate  is  distilled  with  barium  oxide,  under  reduced 
pressure,  hydrazine  is  liberated: 

N2H4,H2O  +  BaO  ->  N2H4|  +  Ba(OH)2. 

Hydrazine  hydrate  freezes  at  about  —  40°  and  boils  at  118.5°. 
Its  aqueous  solution  is  alkaline,  and  salts  are  formed  by  neutraliza- 
tion. 

Hydrazoic  Acid  HN3.  —  When  nitrous  oxide  (q.v.)  is  led  over 
sodamide  at  200°,  water  is  liberated  and  sodium  hydrazoate  remains 

NH2Na  +  N20  ->  NaN3  +  H2O. 

A  dilute  solution  of  the  free  acid  is  best  obtained  by  distilling  the 
lead  salt  with  dilute  sulphuric  acid. 

The  pure  acid  (b.-p.  37°)  is  violently  explosive,  resolving  itself 
into  nitrogen  and  hydrogen  with  liberation  of  much  heat: 

HN3,Aq  -*  H  +  3N  +  Aq  4-  61,600  cal. 

Halogen  Compounds  of  Nitrogen.  —  When  ammonium  chloride 
solution  is  treated  with  excess  of  chlorine,  drops  of  an  oily  liquid, 
nitrogen  trichloride,  are  formed:  3C12  +  NH4C1  ->  NC13  +  4HC1.  It 
is  extremely  explosive,  resolving  itself  into  its  constituents  with 
liberation  of  much  heat. 

When  a  solution  of  iodine  in  potassium  iodide  solution  (p.  165) 
is  added  to  aqueous  ammonia,  a  brown  precipitate  is  formed.  This 
seems  to  have  the  composition  N2H3I3,  and  is  named  nitrogen  iodide. 
It  may  be  handled  while  wet.  When  dry,  if  touched  with  a  feather, 
it  decomposes  into  its  constituents  with  violent  explosion. 

Exercises.  —  1.  When  moist  air  is  used  as  a  source  of  nitrogen, 
what  advantage  is  there  in  using  copper  rather  than  the  less  expen- 
sive metal  iron,  for  removing  the  oxygen  (p.  66)? 

2.  How  many  grams  of  water  at  0°  could  be  frozen  (p.  78)  by  the 
removal  of  the  heat  required  to  evaporate  50  g.  of  liquid  ammonia 
(p.  282)? 

3.  How  many  grams  of  ammonia  are  contained  in  1 1.  of  "  concen- 
trated ammonia  "  (p.  281)? 

4.  What  are  the  ions  of  hydrazine  hydrate?     Formulate  (p.  239) 
the  neutralization  of  this  base  with  sulphuric  acid. 


NITROGEN   AND   AMMONIA  'J86 

5.  What  is  the  object  attained  by  distilling  under  reduced  pres- 
sure in  making  hydrazine  (p.  284)? 

6.  Classify  (pp.  124,  163)  the  interaction  of  a  nitride  with  water 
(p.  281)  and  of  chlorine  and  ammonium  chloride  (p.  282),  and  the 
results  of  heating  ammonium  nitrite  (p.  279)  and  ammonium  chloride 
(p.  283). 


CHAPTER   XXV 
THE    ATMOSPHERE.     THE    HELIUM    FAMILY 

THE  pressure  which  is  exerted  by  the  air  upon  each  square  centi- 
meter of  the  earth's  surface  is  1033.6  g.,  or  a  little  over  one  kilogram. 
This  is  nearly  fifteen  pounds  to  the  square  inch. 

There  are  three  classes  of  components  in  the  air.  Those  of  the  first 
class,  oxygen,  nitrogen,  and  the  inert  gases  of  the  helium  family  (see 
below),  are  present  in  almost  constant  quantities.  Those  of  the  second 
class  are  very  variable  in  quantity,  although  found  in  all  samples  of 
air,  and  include  carbon  dioxide  and  water  vapor.  Those  of  the  third 
class,  such  as  the  sulphur  dioxide  in  city  air,  are  accidental.  Finally, 
one  significant  component  of  the  air  is  the  dust. 

Components  ivhich  are  Constant  in  Amount.  —  In  the  deter- 
mination of  the  oxygen  in  air,  phosphorus  in  the  form  of  thin  wire 
may  be  used.  In  this  way  a  great  surface  is  obtained,  and  the 
absorption  of  oxygen  from  a  measured  sample  of  air  may  be  carried 
out  in  a  few  seconds. 

In  the  air  taken  from  mines,  from  mountain  tops,  from  the  surface 
of  the  sea,  and  from  inland  regions,  the  percentages  of  oxygen  by 
volume  are  found  to  be  fairly  constant,  ranging  between  20.26 
and  21.00,  the  latter  being  the  proportion  in  normal  air. 

When  the  residual  gas  is  led  slowly  through  a  heated  tube  con- 
taining magnesium,  the  nitrogen  unites  with  the  metal  to  form  the 
solid  magnesium  nitride  (p.  280),  and  only  about  10  c.c.  out  of  every 
liter  remains  uncombined.  This  residuum  is  argon,  mixed  with  one- 
hundredth  of  its  volume  of  other  gases  belonging  to  the  helium 
family.  It  is  possible  that  a  trace  of  hydrogen  (cf.  p.  63)  is  one  of 
the  regular  components  of  air. 

Components  ivhich  are  Variable  in  Amount.  —  Pure  country 
air  contains  about  3  parts  in  10,000  of  carbon  dioxide.  In  city  air 
there  are  from  6  to  7  parts  in  the  same  volume,  while  in  the  air  of 
audience-rooms  the  proportion  may  rise  as  high  as  50  parts. 

286 


THE    ATMOSPHEKK  287 

50 

To  determine  the  proportion  of  carbon  dioxide,  a  measured 
volume  of  air  is  bubbled  slowly  through  a  measured  volume  of  a 
solution  of  barium  hydroxide  of  known  concentration.  Barium  car- 
bonate is  precipitated:  Ba(OH)2  -f  CO2  -*  BaCO3 1  +  H2O,  and 
the  quantity  of  barium  hydroxide  remaining  is  determined  by 
titration  (p.  242). 

The  sources  of  the  carbon  dioxide  in  the  air  are  numerous.  It 
comes  from  the  decay  of  vegetable  and  animal  matter,,  in  which, 
chiefly  through  the  influence  of  minute  vegetable  organisms,  the 
carbon  is  oxidized  to  carbon  dioxide.  It  is  formed  also  by  the  com- 
bustion of  coal  and  wood,  and  is  exhaled  by  animals.  The  proportion 
of  this  gas  in  the  air  would  naturally  increase  continuously,  though 
slowly,  as  the  result  of  these  processes,  were  it  not  that  it  is  removed 
just  as  continuously  by  the  action  of  growing  plants  (see  p.  324). 

The  quantity  of  water  vapor  in  the  air  is  constantly  changing.  It 
increases  locally  by  evaporation  from  the  soil  and  from  natural  waters, 
particularly  in  warm  weather.  It  decreases  when  local  cooling  leads 
to  the  precipitation  of  water  in  the  forms  of  mist  and  rain.  The 
phrase  commonly  heard,  that  on  a  moist  day  the  atmosphere  is 
"  laden  "  with  moisture,  is  peculiarly  inapt.  We  recognize  at  once 
from  observation  of  the  barometer,  which  is  lower  in  such  a  state  of 
the  atmosphere,  that  the  pressure  of  the  air  is  less.  Moist  air  must 
be  lighter  than  dry  air,  for  in  it  a  certain  proportion  of  water  mole- 
cules, of  molecular  weight  18,  is  substituted  for  an  equal  number 
(cf.  p.  125)  of  molecules  of  nitrogen  and  oxygen  whose  relative 
weights  are  28  and  32  respectively.  The  result  is  therefore  a  diminu- 
tion in  the  specific  gravity  of  the  air.  The  proportion  of  water  in  a 
given  volume  of  air  may  be  measured  most  accurately  by  permitting 
the  air  to  stream  slowly  through  tubes  filled  with  calcium  chloride 
or  phosphoric  anhydride.  The  increase  in  weight  of  the  charged 
tubes  represents  the  quantity  of  moisture  abstracted  from  the 
sample. 

The  dust  varies  both  in  kind  and  quantity  according  to  the  locality. 
It  is  found  to  be  partly  inorganic.  The  organic  dust  may  be  divided 
into  two  kinds.  The  part  which  is  dead  includes  coal  dust,  refuse 
from  the  streets,  minute  shreds  of  cotton,  linen,  hay,  etc.  The 
living  dust  consists  of  pollen  grains,  spores  of  fungi  and  other  plants, 
bacteria,  and  similar  microscopic  organisms.  The  presence  of  micro- 
scopic germs  in  the  air  is  shown  by  the  fact  that  when  nutritive 


288  COLLEGE   CHEMISTRY 

liquids  have  been  exposed  to  the  air,  even  for  a  few  minutes,  putre- 
faction very  soon  sets  in.  Some  germs  also  produce  disease  when 
they  gain  access  to  the  body,  particularly  through  wounds,  or  inci- 
sions made  in  the  course  of  operations.  The  object  of  antiseptic 
treatment  by  the  use  of  phenol  (carbolic  acid),  mercuric  chloride, 
and  other  substances,  is  to  destroy  such  organisms  or  to  hinder  their 
development. 

Flasks  can  be  filled  with  dustless  air  through  the  displacement  of 
that  which  they  contain  by  air  drawn  through  a  wide  tube  packed 
with  12-15  inches  of  cotton.  It  has  been  shown  by  Ait  ken  that 
dustless  air  behaves  differently  from  ordinary  air  in  respect  to  the 
way  in  which  its  moisture  condenses.  In  ordinary  air,  the  dust 
particles  act  as  nuclei,  and  a  fog  is  formed.  In  dustless  air,  no  fog 
can  be  produced. 

Air  a  Mixture.  —  Since  the  main  components  of  air  were  not 
definitely  identified  until  the  end  of  the  eighteenth  century,  we  can 
understand  why  the  substance  was  for  long  considered  to  be  an 
element.  The  experiments  which  we  have  described,  in  which  the 
oxygen  was  removed  from  the  air  and  the  nitrogen  remained,  do  not 
prove  that  the  original  constituents  were  present  simply  in  mechani- 
cal mixture.  They  might  have  been  combined,  and  the  combustion 
of  phosphorus,  for  example,  might  have  represented  the  removal  of 
oxygen  from  combination  with  nitrogen  and  its  appropriation  by 
the  phosphorus.  It  may  be  well,  therefore,  to  point  out  some 
reasons  which  lead  us  to  regard  the  air  as  a  mixture: 

1.  When  oxygen  and  nitrogen  are  mixed  in  the  proper  proportions, 
we  obtain  a  gas  identical  with  air  in  all  its  properties,  and  there  is  no 
evidence  of  any  production  or  absorption  of  heat,  such  as  would 
occur  in  case  of  chemical  combination. 

2.  The  proportion  by  volume  is  not  perfectly  constant. 

3.  The  proportions  by  weight   in   which   the   components  are 
contained  in  air  are  not  integral  multiples  of  the  atomic  weights. 

Composition  of  Air.  —  Air,  when  freed  from  carbon  dioxide  and 
water,  contains  by  volume  78.06  per  cent  of  nitrogen,  21.00  per  cent 
of  oxygen,  and  0.94  per  cent  of  argon.  When  only  the  water  is 
removed,  the  carbon  dioxide  averages  about  0.03  per  cent  of  the 
whole. 


THE    ATMOSPHERE  289 

Graham  suggested  an  illustration  which  will  make  those  propor- 
tions clearer.  He  says  that  if  we  imagine  the  air  to  be  divided  by 
magic  into  its  components,  and  to  remain  separated  for  a  time  suffi- 
ciently long  to  enable  us  to  note  the  proportions,  and  if  the  substances 
arrange  themselves  in  the  order  of  their  specific  gravities,  we  should 
have  the  following  layers  resting  upon  the  surface  of  the  earth  and 
one  upon  another:  On  the  earth,  five  inches  of  water;  above  that, 
thirteen  feet  of  carbon  dioxide;  above  that,  a  mile  of  oxygen;  and  on 
the  top,  about  four  miles  of  nitrogen.  This  would  be  on  the  assump- 
tion that  these  gases  were  compressed  so  as  to  have  the  same  density 
throughout.  We  should  now  add  a  layer  of  argon,  of  about  ninety 
yards  thickness,  between  the  carbon  dioxide  and  oxygen. 

Air  and  Health.  —  As  human  beings  we  have  an  especial  interest 
in  the  composition  of  the  air,  since  our  living  depends  upon  the 
oxygen  which  we  secure  by  breathing  it  (cf.  p.  52);  We  draw  about 
half  a  liter  of  air  into  our  lungs  at  each  breath,  or  about  half  a  cubic 
meter  per  hour.  The  oxygen  of  this  air  is  partly  used,  being  taken 
up  by  the  blood,  and  part  remains  in  the  exhaled  air.  On  the  other 
hand,  carbon  dioxide  is  given  off  in  the  lungs  and  passes  out  with 
the  unused  oxygen.  The  nitrogen  is  unaffected.  In  100  c.c.  of 
expired  air  there  are  contained  about  15.9  c.c.  of  oxygen  and  4.5  c.c. 
of  carbon  dioxide.  The  total  quantity  of  oxygen  consumed  during 
twenty-four  hours  is  about  three-fourths  of  a  kilogram,  or  more  than 
half  a  cubic  meter.  While  the  greater  part  of  this  gains  access  to 
the  body  through  the  lungs,  more  or  less  exchange  of  gases  takes 
place  in  all  animals  through  the  skin.  The  lower  limit  of  oxygen  for 
respirable  air  is  about  10  per  cent,  although  a  candle  goes  out  when 
the  proportion  reaches  about  18.5  per  cent. 

Liquefaction  of  Gases.  —  The  earliest  experiments  of  this  kind 
were  made  by  Northmore  (1805),  who  liquefied  chlorine,  hydrogen 
chloride,  and  sulphur  dioxide.  In  1823  chlorine  was  again  liquefied 
by  Faraday.  During  the  following  years  he  reduced  sulphur  dioxide, 
hydrogen  sulphide,  carbon  dioxide,  nitrous  oxide,  cyanogen,  and 
ammonia  to  the  liquid  condition.  In  1883  Wroblewski  and  Olszew- 
sld  prepared  visible  amounts  of  liquid  oxygen.  About  the  same 
time  Dewar  devised  means  of  manufacturing  large  quantities  of 
liquid  air  and  oxygen.  The  most  successful  apparatus  for  use  on 
a  small  scale  is  that  devised  by  Hampson. 


290  COLLEGE    CHEMIST.RY 

Liquid  Air.  —  Liquid  air  varies  in  composition,  as  the  nitrogen 
(b.-p.  —  194°)  is  less  condensible  than  the  oxygen  (b.-p.  -  182.5°). 
It  boils  at  about  —  190°,  and  contains  about  54  per  cent  of  oxygen 
by  weight,  while  air  contains  23.2  per  cent.  By  allowing  evapora- 
tion to  go  on,  a  liquid  containing  75  to  95  per  cent  of  oxygen  is  easily 
obtained  (cf.  p.  46).  The  gas  secured  by  the  evaporation  of  the 
residue  is  pumped  into  cylinders  and  sold  as  compressed  oxygen. 
Cartridges  made  of  granular  charcoal  and  cotton  waste,  when  satu- 
rated with  liquid  air,  have  been  used  as  an  explosive  in  mining. 

THE  HELIUM  FAMILY. 

Argon.  —  Lord  Rayleigh  was  the  first  to  observe  that,  while 
specimens  of  oxygen  and  other  gases  made  purposely  from  various 
sources  always  had  the  same  density,  nitrogen  was  an  exception. 
One  liter  of  nitrogen  made  from  air,  and  supposed  to  be  pure, 
weighed  1.2572  g.  When  the  gas  was  manufactured  by  decompo- 
sition of  five  different  compounds,  such  as  urea  and  certain  oxides  of 
nitrogen,  the  results  agreed  well  amongst  themselves.  The  mean 
weight  of  a  liter  of  this  nitrogen  was  only  1.2505  g.  The  difference, 
amounting  to  nearly  7  mg.,  was  very  much  greater  than  the  experi- 
mental error.  The  suspicion  naturally  arose  that  some  heavier  gas 
was  present  in  natural  nitrogen.  Soon  after  (1894),  Professor,  now 
Sir  William  Ramsay  obtained  argon  by  removal  of  the  greatly  pre- 
ponderating nitrogen  by  means  of  magnesium  (p.  280).  The  new 
gas  had  a  molecular  weight  of  about  40,  and  was  therefore  more 
than  one-third  heavier  than  nitrogen. 

The  exact  density  of  argon  is  39.9.  When  liquefied  it  boils  at 
—  186°,  and  the  colorless  solid  melts  at  —  189.5°.  The  solubility 
of  the  gas  in  water  (4  volumes  in  100)  is  two  and  one-half  times  that 
of  nitrogen.  It  has  not  been  found  to  enter  into  any  sort  of  chemical 
combination,  and  was  named  argon  on  this  account  (Gk.  d/oyos, 
inactive).  The  physical  properties  show  that  the  molecules  of  the 
gas,  like  those  of  mercury  (p.  138),  are  monatomic. 

Helium.  —  In  1868  Lockyer  first  detected  an  orange  line  in  the 
spectrum  of  the  sun's  prominences  which  was  not  given  by  any 
terrestrial  substance  then  known.  The  line  was  so  conspicuous  that 
it  was  attributed  to  the  presence  of  a  new  chemical  element,  which 
was  named  helium  (Gk.  17X109,  the  sun).  Ramsay,  in  searching  for 


THE    HELIUM    FAMILY  291 

sources  of  argon,  examined  the  "  nitrogen  "  which  was  reported  by 
various  mineralogists  as  being  disengaged  when  certain  rare  minerals 
were  heated.  These  minerals,  cleveite,  uraninite,  and  broggerite, 
were  chiefly  compounds  of  uranium,  yttrium,  and  thorium.  He  was 
surprised  to  find  (1895)  that  the  gas  was  not  always  nitrogen.  It 
frequently  contained  a  large  proportion  of  a  very  light  gas,  the  spec- 
trum of  which  was  identical  with  that  of  solar  helium.  The  same 
gas  is  found  in  small  amount  in  the  atmosphere.  Helium  does  not 
exhibit  any  tendency  to  enter  into  combination.  It  is  monatomic 
and  its  density  shows  that  its  molecular  weight  is  4.  It  has 
neither  been  liquefied  nor  solidified. 

Neon,  Krypton,  and  Xenon.  —  When  the  argon  obtained  from 
atmospheric  nitrogen  is  cooled  with  liquid  air  (—  185°),  the  argon, 
krypton,  and  xenon  are  liquefied,  and  the  neon  and  helium  are  dis- 
solved by  the  liquid.  When  heat  is  allowed  to  reach  the  mixture, 
the  last  two  gases  escape  first,  along  with  much  argon.  When  most 
of  the  argon  has  escaped,  the  krypton  and  xenon  still  remain  liquid. 
By  repeated  liquefaction  and  fractional  evaporation  (see  under 
Petroleum),  the  krypton  and  xenon  are  separated  from  the  argon 
and  from  one  another.  When  the  vessel  containing  the  mixture  of 
helium  and  neon  is  immersed  in  liquid  hydrogen  (—  240°),  the  second 
freezes  to  a  white  solid,  and  the  helium,  which  remains  gaseous,  can 
be  pumped  off. 

These  gases  are  all  entirely  inactive  chemically,  and  are  all  mon- 
atomic. Their  molecular  weights  are:  Neon,  20;  krypton,  81.5; 
xenon,  128. 

Exercises. —  1.  A  sample  of  moist  air,  confined  over  water  at 
15°  and  760  mm.,  occupies  15  c.c.  It  is  mixed  with  20  c.c.  of  hydro- 
gen, and  the  mixture  is  exploded,  and  suffers  a  contraction  of  9.5  c.c. 
What  would  be  the  volume  of  the  oxygen  it  contained  if  measured 
dry  at  0°  and  760  mm.? 

2.  Calculate,  from  the  data  on  p.  286  and  the  densities,  the 
percentage  by  weight  of  the  three  principal  components  of  air. 

3.  Of  the  proofs  that  air  is  a  mixture  (p.  288),  which  shows  that 
no  part  of  the  components  is  combined,  and  which  that  the  com- 
ponents are  not  wholly  combined? 


CHAPTER  XXVI 
OXIDES  AND  OXYGEN  ACIDS  OF  NITROGEN 

THE  names  and  formulas  of  the  oxides  and  oxygen  acids  of  nitrogen 
are  as  follows: 

Nitrous  oxide  N2O  < Hyponitrous  acid  H2N2O2 

Nitric  oxide  NO 

Nitrous  anhydride  N2O3         < — >       Nitrous  acid  HN02 

Nitrogen  tetroxide  N2O4  and  N02 

Nitric  anhydride  N2O5  < — >      Nitric  acid  HNO3. 

All  the  oxides  are  endo thermal  compounds  (p.  55),  yet,  with  the 
exceptions  of  the  third  and  the  last,  they  are  all  relatively  stable. 
The  acids,  when  deprived  of  the  elements  of  water,  yield  the  oxides 
opposite  which  they  stand.  Conversely,  excepting  in  the  case  of 
nitrous  oxide,  the  anhydrides  with  water  give  the  acids.  All  of 
these  substances  are  obtained  directly  or  indirectly  from  nitric  acid  — 
nitric  anhydride  by  removal  of  water,  the  others  by  reduction.  We 
turn,  therefore,  first,  to  nitric  acid,  its  sources  and  properties. 

NITRIC   ACID   HN03. 

Sources.  —  Sodium  nitrate,  or  Chili  saltpeter,  is  found  in  a  desert 
region  near  the  boundary  of  Chili  and  Peru.  The  deposit  is  about 
5  feet  thick,  2  miles  wide,  220  miles  long,  and  contains  20  to  55  per 
cent  of  the  salt.  Purification  is  effected  by  recrystallization.  Potas- 
sium nitrate,  or  Bengal  saltpeter,  is  found  in  the  soil  in  the  neighbor- 
hood of  cities  in  India,  Persia,  and  other  oriental  countries.  It  arises 
from  the  oxidation  of  animal  refuse  through  the  mediation  of  nitri- 
fying bacteria.  The  potash  and  lime  in  the  soil,  along  with  the 
product  of  oxidation  of  the  nitrogen,  give  nitrates  of  potassium  and 
calcium.  The  aqueous  extract  of  this  soil  is  treated  with  wood  ashes, 
on  account  of  the  potash  (K2C03)  contained  in  them.  It  is  poured 
off  from  the  calcium  carbonate  thus  precipitated,  and  is  finally 
evaporated. 

292 


OXIDES   AND   OXYGEN   ACIDS   OF   NITROGEN  293 

Preparation.  —  When  any  nitrate  is  treated  with  any  acid,  nitric 
acid  is  formed  by  a  reversible  double  decomposition.  As  sodium 
nitrate  is  the  cheapest  salt  of  nitric  acid,  it  is  always  employed.  For 
the  same  reason  of  low  cost,  and,  above  all,  because  of  its  relative 
in  volatility,  sulphuric  acid  is  used  to  displace  it: 

NaNO3  +  H2SO4^NaHSO4  +  HN03t. 

The  nitric  acid  is  rather  volatile  (b.-p.  86°).  while  sulphuric  acid  (b.-p. 
330°)  is  much  less  so,  and  the  two  salts  are  not  volatile  at  all.  Thus 
the  interaction  proceeds  to  completion  very  easily  (cf.  p.  184).  The 
materials  are  heated  in  cast-iron  stills,  and  the  vapor  is  condensed  in 
earthenware  pipes  surrounded  by  water. 

Another  action  by  which  attempts  are  being  made  to  manufacture 
nitric  acid  is  the  direct  union  of  the  nitrogen  and  oxygen  of  the  air 
under  the  influence  of  an  electric  discharge.  The  nitrogen  tetroxide 
(NO2),  which  is  formed  in  small  amounts  at  a  time,  is  dissolved  in 
water:  WQ  +  J^Q  _,  2HNO  +  NO. 


The  nitric  oxide  gas,  on  escaping  from  the  water,  unites  directly  with 
oxygen  to  reproduce  the  tetroxide.  The  reaction  is  of  interest,  inde- 
pendently of  this  one  application,  because  of  its  reversibility.  It  pro- 
ceeds forward  with  excess  of  water,  while  the  reverse  action  takes 
place  when  nitric  oxide  (q.v.)  comes  in  contact  with  concentrated 
nitric  acid  in  which  the  quantity  of  water  is  at  a  minimum. 

Physical  Properties.  —  Nitric  acid  is  a  colorless,  mobile  liquid, 
boiling  at  86°,  and  freezing  to  a  solid  which  melts  at  -  47°.  It 
fumes  strongly  when  its  vapor  issues  into  moist  air  (cf.  p.  120).  An 
aqueous  solution  containing  68  per  cent  of  the  acid  boils  at  120.5°, 
while  the  pure  acid,  pure  water,  and  all  other  mixtures,  boil  at  lower 
temperatures.  This  68  per  cent  nitric  acid  of  constant  boiling-point 

(p.  120)  forma  the  "  concentrated  nitric  acid  "  of  commerce. 

i 

Chemical  Properties.  —  1.  Like  chloric  acid  (p.  196),  and  other 
oxygen  acids  of  the  halogens,  nitric  acid  is  most  stable  when  mixed 
with  water.  The  pure  (100  per  cent)  acid  decomposes  while  being 
distilled:  4HN03  ->  4NO2  +  2H2O  +O2, 

yet  not  with  explosive  violence  like  chloric  acid.  The  distillate  is 
colored  brown  by  dissolved  nitrogen  tetroxide  (N02).  Repeated 


294  COLLEGE    CHEMISTRY 

distillation  finally  leaves  68  per  cent  of  the  acid,  mixed  with  32  per 
cent  of  water  formed  by  the  above  decomposition.  The  acid  of  con- 
stant boiling-point  is,  therefore,  reached,  as  usual,  from  more  concen- 
trated as  well  as  from  less  concentrated  specimens. 

"  Fuming  "  nitric  acid  is  brown  in  color,  and  contains  a  consider- 
able amount  of  dissolved  nitrogen  tetroxide.  It  is  made  by  distilling 
the  acid  with  a  little  starch.  The  latter  reduces  a  part  of  the  nitric 
acid  and  liberates  more  of  the  tetroxide  than  does  mere  distillation. 

2.  Nitric  acid,  when  dissolved  in  water,  is  highly  ionized,  and 
is  therefore  active  as  an  acid.     By  interaction  with  hydroxides  and 
oxides  it  forms  nitrates. 

3.  When  pure  nitric  acid  (b.-p.  86°)  is  poured  upon  phosphoric 
anhydride,  the  latter  combines  with  the  elements  of  water,  and 
distillation  gives  nitric  anhydride:  2HN03  +  P2O5  ->  N2O5  f  -f  2HP03. 
The  anhydride  is  a  white  solid  melting  at  30°  and  boiling  at  45°.     It 
unites  vigorously  with  water  to  form  nitric  acid.     It  decomposes 
spontaneously  into  nitrogen  tetroxide  and  oxygen,  2N2O5— >4N02  +  02. 

4.  Like  the  unstable  oxygen  acids  of  the  halogens,  nitric  acid  is  an 
oxidizing  agent  even  when  diluted  with  water.     The  multiplicity  of 
the  products  into  which  it  may  be  decomposed  by  reduction,  however, 
renders  separate  treatment  of  this  property  necessary  (see  p.  297). 

5.  Nitric  acid  interacts  energetically  with  many  compounds  of 
carbon  to  give  nitro-derivatives.     Thus,  when  heated  with  phenol 
(carbolic  acid)  it  gives  picric  acid  (trinitrophenol,  HC0H2(NO2)3O), 
which  crystallizes  in  yellow  needles  in  the  mixture: 

C6H5(OH)  +  3HONO2-»C6H2(OH)(NO2)3  +  3H2O. 

6.  Organic  compounds  of  another  class,  the  alcohols  (q.v.\  inter- 
act with  molecular"  nitric  acid  in  a  different  way.     The  latter  i& 
mixed  with  sulphuric  acid,  which  assists  in  the  removal  of  the  ele- 
ments of  water  (p.  263).     Thus,  when  glycerine  is  added  slowly  to 
the  cooled  mixture,  glyceryl  nitrate   (so-called  nitro-glycerine)   is 
produced: 

C3H5(OH)3  +  3HN03-»C3H5(N03%  +  3H20. 

Gun-cotton  is  made  by  this  action,  cotton  (cellulose)  being  employed : 

S(C6H1005)2  +  6HN03-*C12H1404(N03)6  +  6H2O. 

7.  Nitric,  acid  produces  substances  of  bright-yellow  color,  known 
as  xanthoproteic  acids,  when  it  comes  in  contact  with  the  skin. 


OXIDES   AND   OXYGEN   ACIDS   OF   NITROGEN  295 

Nitrates.  -^:.The  nitrates  are  all  more  or  less  easily  soluble  in 
water.  When  heated  they  decompose  in  one  or  other  of  three  ways 
(see  pp.  296,  299,  300).  The  individual  nitrates,  such  as  sodium 
nitrate  and  potassium  nitrate,  are  described  elsewhere. 

NITRIC  OXIDE  AND  NITROGEN  TETROXIDE. 

Preparation  of  Nitric  Oxide  NO.  —  Pure  nitric  oxide  is  ob- 
tained^by  adding  nitric  acid  to  a  boiling  solution  of  ferrous  sulphate 
in  dilute'stilphuric  acid  or  of  ferrous  chloride  in  hydrochloric  acid: 

2FeS04  +  H2SO4  -» Fe,(S04)8  (+  2H)    x  3          (1) 

(3H)  +  HN03  -»  NO  +  2H2O  x  2          (2) 

6FeSO4  +  3H2SO4  +  2HNO3-»3Fe2(SO4)3  +  2NO  +  4H2O. 

The  first  partial  equation  does  not  take  place  at  all  unless  an  oxidiz- 
ing agent  like  nitric  acid  is  present  (p.  213).  The  multiplication  of 
the  two  partial  equations  by  3  and  2  respectively  is  required  in  order 
that  the  hydrogen,  which  is  not  a  product,  may  cancel  out.  This 
action  may  be  used  as  a  means  of  determining  the  quantity  of 
nitric  acid  in  a  solution,  or  of  nitrates  in  a  mixture,  by  measure- 
ment of  the  volume  of  nitric  oxide  evolved. 

As  we  shall  see,  this  gas  may  also  be  obtained  when  sufficiently 
dilute  nitric  acid  (sp.  gr.  1.2^  acts  upon  copper.  Although  some 
nitrous  oxide  and  nitrogen  are  produced  in  this  interaction,  it  fur- 
nishes a  convenient  method  of  generating  the  gas. 

Properties  of  Nitric  Oxide.  —  Nitric  oxide  is  a  colorless  gas. 
In  solid  form  it  melts  at  -  150°,  and  the  liquid  boils  at  —  142.4° 
under  757.2  mm.  pressure.  Its  solubility  in  water  is  slight. 

The  density  of  the  gas  shows  the  formula  to  be  NO;  and  there  is  no 
tendency  to  form  a  polymer,  such  as  N202,  even  at  low  temperatures. 
This  gas  is  the  most  stable  of  the  oxides  of  nitrogen.  Vigorously 
burning  phosphorus  continues  to  burn  in  the  gas.  Burning  sulphur 
and  an  ignited  taper,  however,  are  extinguished. 

Nitric  oxide  has  two  characteristic  properties.  It  unites  directly 
with  oxygen  in  the  cold  to  form  the  reddish-brown  nitrogen  tetroxide: 

2NO  +  O2  <=>  2N02. 

The  same  result  follows  when  it  is  led  into  warm  concentrated  nitrio 
acid  (cf.  p.  293) :  NO  +  2HN08  ?±  3N02  +  H20. 


296  COLLEGE    CHEMISTRY 

It  also  unites  with  a  number  of  salts,  the  compound  in  the  case  of 
ferrous  sulphate,  FeS04,NO,  being  capable  of  existence  in  solution 
and  possessing  a  brown  color.  Since  ferrous  sulphate  will  first 
reduce  nitric  acid  to  nitric  oxide  (p.  295),  and  the  excess  of  the  salt 
will  then  give  a  brown  color  with  the  product,  a  delicate  test  for 
nitric  acid  is  founded  upon  these  actions. 

Preparation  of  Nitrogen  Tetroxide  NO2. —  This  substance 
is  liberated  by  heating  nitrates,  other  than  those  of  potassium, 
sodium,  or  ammonium: 

2Cu(NO3)2-»2CuO  +  4NO2  +  O2. 

In  most  cases  the  oxide  of  the  metal  remains.  When  the  mixed 
gases  are  led  through  a  U-tube  immersed  in  a  freezing  mixture,  the 
tetroxide  condenses  as  a  pale-yellow  liquid  (b.-p.  22°),  and  the  oxygen 
passes  on. 

The  compound  may  also  be  made  by  direct  union  of  nitric  oxide 
and  oxygen,  or  by  oxidation  of  nitric  oxide  by  concentrated  nitric 
acid  (p.  295).  It  is  likewise  almost  the  sole  product  of  the  inter- 
action of  concentra£dt nitric  acid  and  tin  or  copper  (see  p.  298).  If  any 
nitric  oxide  were  produced  by  the  primary  action,  it  would  be  oxidized 
to  nitrogen  tetroxide  in  passing  up  through  the  acid  (p.  295). 

Properties  of  Nitrogen  Tetroxide.  —  The  most  striking 
peculiarity  of  this  gas  is  that,  when  hot,  it  is  deep  brown  in  color,  and 
when  cold,  pale  yellow.  The  density  of  the  brown  gas,  at  140°, 
corresponds  to  the  formula  NO2,  that  of  the  yellow  gas  at  22°,  to 
N2O4.  When  the  temperature  is  carried  above  154°,  by  passing  the 
brown  gas  through  a  red-hot  tube,  the  brown  color  disappears,  and 
nitric  oxide  and  oxygen  are  formed.  On  cooling,  the  same  steps 
through  brown  gas  to  pale-yellow  gas  are  retraced: 

2NO  +  O2  *±  2NO2  *±  N304, 
Colorless  Brown      Colorless 

Since  nitrogen  tetroxide  yields  free  oxygen  more  readily  than  does 
nitric  oxide,  phosphorus  burns  readily  in  it:  a  taper,  however,  is 
extinguished.  It  has  powerful  oxidizing  properties;  and  " fuming  " 
nitric  acid  (p.  294),  which  contains  it  in  solution,  is  employed  when 
oxidation  is  the  special  object  in  view,- 


OXIDES  AND   OXYGEN   ACIDS   OF   NITROGEN  297 

This  oxide  is  intermediate  in  composition  between  nitrous  and 
nitric  anhydrides,  and,  when  dissolved  in  cold  water,  gives  both 
nitric  and  nitrous  acids:  N2O4  +  H2O  ->  HNO3  -f  HN02.  If  a  base 
is  present,  a  mixture  of  the  nitrate  and  nitrite  of  the  metal  is  produced. 
When  the  water  is  not  cooled,  the  nitrous  acid  (q.v.),  being  unstable, 
gives  nitric  oxide  and  nitric  acid,  so  that  the  result  is:  3N03  +  H2O 
*=±  2HNO3  +  NO. 

OXIDIZING  ACTIONS  OP  NITRIC  ACID. 

When  nitric  acid  gives  up  oxygen  to  any  body,  it  is  itself  reduced. 
Hence,  according  to  convenience,  we  shall  refer  to  oxidations  by,  or 
reductions  of  nitric  acid. 

Oxidation  of  Hydrogen.  —  The  metals  preceding  hydrogen 
in  the  electromotive  series  (p.  245),  displace  hydrogen  from  nitric 
acid,  as  they  do  from  other  acids.  With  metals  more  active  than 
zinc,  such  as  magnesium,  a  great  part  of  the  hydrogen  escapes  in  the 
free  condition.  But,  in  the  case  of  zinc  and  the  metals  below  it, 
most  or  all  of  the  hydrogen  is  oxidized  to  water  by  the  nitric  acid,  and 
part  of  the  acid  is  reduced  (see  Active  hydrogen,  p.  302).  Thus, 
with  zinc  and  very  dilute  nitric  acid,  almost  the  only  product,  aside 
from  zinc  nitrate,  is  ammonia: 


4Zn  4 

-    8HN03 

—  > 

4Zn(N03)2  (4 

•8H) 

(8HH 

HN03 

—  » 

NH3  +  3H2O 

NHj-4 

HNO3 

-* 

NH4NO3 

4Zn  H 

-  10HNO3 

—  > 

4Zn(NO3)2  + 

NH4NO3  H 

-  3H2O. 

(1) 
(2)  V 

(3) 


With  the  excess  of  nitric  acid  (3),  ammonium  nitrate  is  formed. 


Hecwy  Metal*.  —  The  lees  active  metals,  such  as  copper  and 
silver,  do  not  displace  hydrogen  from  dilute  acids  (p.  245),  but 
reduce  nitric  acid,  nevertheless,  and  are  converted  into  nitrates. 
Platinum  and  gold  (cf.  p.  264)  alone  are  not  attacked.  Thus,  copper, 
with  somewhat  diluted  nitric  acid,  gives  cupric  nitrate  and  nitric 
oxide  (NO). 

In  making  the  equation  for  this  action  we  may  use  a  plan  which  is 
applicable  whenever  an  oxygen  acid  gives  an  oxide  by  reduction 
(cf.  p.  260).  We  resolve  the  formula  of  nitric  acid  into  those  of 


298  COLLEGE    CHEMISTRY 

water  and  the  anhydride  H20,N2O5  (=  2HNO3).  This  shows  that 
the  two  molecules  of  the  acid  will  give  2NO,  and  30  will  remain: 

2HN03-»H,0    +2NO(+30)  (1) 

(30)  +  6HNO3  -f  3Cu->3H2O  +  3Cu(NO3)2  (2) 

8HN03  +  3Cu  -»  4H,O  +  2NO  +  3Cu(NO3)2. 

The  nitric  oxide  is  liberated  as  a  colorless  gas,  but  forms  the  brown 
tetroxide  at  once  on  meeting  the  oxygen  of  the  air  (p.  295). 

When  concentrated  nitric  acid  is  used  with  copper,  almost  pure 
nitrogen  tetroxide  is  obtained: 

2HNO3  -»  H2O  +  2N02  (+  0)  (1) 

(0)  +  2HN03  +  Cu  ->  H2O  +  Cu(N03)2 (2) 

4HNO3  +  Cu->2H2O+   2NO2  +  Cu(NO3)2. 

The  reader  should  note  the  constant  production  of  nitric  oxide  with 
diluted  nitric  acid,  and  the  invariable  formation  of  nitrogen  tetroxide 
with  concentrated  acid.  This  is  explained  by  the  fact  that  nitrogen 
tetroxide  cannot  pass  unchanged  through  a  liquid  containing  much 
water,  for  it  gives  nitric  acid  and  nitric  oxide  with  the  latter  (p.  297). 
Conversely,  where  the  nitric  acid  is  concentrated,  nitric  oxide,  even 
if  formed  by  the  interaction  with  the  metal,  must  be  oxidized  to 
nitrogen  tetroxide  as  it  passes  up  through  the  liquid  (p.  295). 

Oxidation  of  Non-Metals.  —  With  non-metals  the  actions  are 
different,  in  so  far  that  these  elements  do  not  give  nitrates.  Thus 
sulphur  boiled  in  nitric  acid  gives  sulphuric  acid,  along  with  nitric 
oxide,  equation  (3),  or  with  nitrogen  tetroxide,  equation  (6),vor  with 
both,  according  to  the  concentration  of  the  acid  (see  above^ 

2HNO,  U  2NO  +  H2O  (+  30)  (1) 

(3O)  +  H,O  +  S  -»  H2S04 (2) 

2HNO3  +  S  -»  2NO  +  H2SO4  (3) 

2HN03  -»  2N02  +  H2Q    +  O      X  3  (4) 

(3O)  +  H,O  +  S  -»  H2SO4 (5) 

6HNO3  +  S  ->  6NO2  +  2H2O  +  H2S04  (6) 

The  reader  will  note  (cf.  p.  194)  that  a  separate  equation,  (3)  and  (6), 
must  be  made  for  the  formation  of  each  reduction  product.  If  NO 
and  N02  are  both  formed,  they  cannot  arise  from  the  same  molecule 
of  nitric  acid.  They  result  from  two  actions  which  are  independent, 


OXIDES   AND    OXYGEN   ACIDS   OF   NITROGEN  299 

although  proceeding  simultaneously  in  the  same  vessel  (c/.  p.  196). 
Thus  the  equation:  2HN03  +  C  ->  H20  +  CO2  +  ]TO  +  NO2,  is  a 
misrepresentation.  It  implies  that  equimolar  quantnres  of  the  two 
oxides  of  nitrogen  are  formed.  But  this  could  only  occur  by  chance, 
and  the  balance  would  be  destroyed  the  next  moment  by  the  lower- 
ing in  the  concentration  of  the  acid,  giving  the  advantage  to  the  nitric 
oxide. 

Oxidation  of  Compounds :  Aqua  Regia.  -—  Compounds  like 
hydrogen  sulphide  and  sulphurous  acid,  wtffch  are  easily  oxidized, 
interact  with  nitric  acid.  With  diluted  nitric  acid,  the  products  are 
free  sulphur  and  sulphuric  acid  respectively. 

The  mixture  of  nitric  acid  and  hydrochloric  acid  is  known  as  aqua 
regia.  The  chlorine  set  free  by  the  oxidation  of  the  hydrochloric  acid 
is  more  active  than  is  the  ordinary  solution  of  chlorine  in  water,  per- 
haps in  consequence  of  catalytic  action  of  the  substances  in  this 
solution  (see  Active  hydrogen,  p.  302),  and  combines  with  gold  and 
platinum  (q.v.),  converting  them  into  chlorides.  Nitrosyl  chloride 
(NOC1),  which,  however,  does  not  interact  directly  with  the  noble 
metals,  is  formed  also:  « 

O 


C12  +  a  -  N  -  O 

The  interaction  with  platinum  gives  chloroplatinic  acid: 
2HC1  +  2C12  +  Pt  ->  H2PtCl6. 

NITROUS  ACID,  HYPONITROUS  ACID,  AND  THEIR  ANHYDRIDES. 

Nitrites  and  Nitrous  Acid.  —  When  the  nitrates  of  potassium 
and  sodium  are  heated,  they  lose  one  unit  of  oxygen,  and  the 
nitrites  remain: 

2NaNO3  ->  2NaNO2  +  O2. 

Commonly  lead  is  stirred  with  the  melted  nitrate  and  assists  in  the 

removal  of  the  oxygen.     The  litharge  (PbO)  which  is  formed  remains 

as  a  residue  when  the  sodium  nitrite  is  dissolved  for  recrystallization. 

When  an  acid  is  added  to  a  dilute  solution  of  a  nitrite,  a  pale-blue 

solution  containing  nitrous  acid  HNO2  is  obtained.     The  acid  is  very 

unstable,  however,  and,  when  the  solution  is  warmed,  it  decomposes: 

3HN02  J*  HN03  4-  2NO  +  H,O. 


300  COLLEGE   CHEMISTRY 

When  a  concentrated  solution  of  sodium  nitrite  is  acidified,  the 
nitrous  acid  decomposes  at  once,  and  a  brown  gas  containing  the 
anhydride  escapes: 

2H*  +  2NO/  s=?  2HNO2  fc?  H2O  +  N2O3 1. 

This  behavior  distinguishes  a  nitrite  from  a  nitrate. 

Reducing  agents  deprive  nitrous  acid  of  part  or  all  of  its  oxygen: 
2HI  +  2HNO2  ->  2H2O  +  2NO  +  I2. 

Indigo  is  also  converted  by  it  into  isatin  (cf.  p.  192).  On  the  other 
hand,  oxidizing  agents  which  are  sufficiently  active,  like  acidified 
potassium  permanganate,  convert  nitrous  acid  into  nitric  acid: 

3H2SO4 + 2KMnO4^K2SO4  +  2MnSO4 + 3H2O  ( + 50)    (1) 

(5O)  +  5HNO2  -»  5HNO3 (2) 

3H2SO4 + 2KMnO4 + 5HNO2-»K3SO4  -f  2MnSO4 + 3H2O  +  5HNO3 

Nitrous  acid  is  much  used  in  the  making  of  organic  dyes. 

Nitrous  Anhydride  N2O3.  —  A  study  of  the  gas  arising  from 
the  decomposition  of  nitrous  acid  shows  that  in  the  gaseous  state 
the  anhydride  is  almost  entirely  dissociated: 

N2O3  <=>  NO  +  NO2. 

When  the  mixture  is  led  through  a  U-tube  immersed  in  a  freezing 
mixture  at  —  21°,  a  deep-blue  liquid  is  obtained  which  is  the 
anhydride  itself.  This  dissociates  when  allowed  to  boil. 

The  same  equimolar  mixture  of  the  two  gases  is  obtained  by  the 
action  of  water  on  nitrosylsulphuric  acid  (p.  259). 

Hyponitrous  Acid  and  Nitrous  Oxide  N2O.  —  Hyponitrous 
acid  H2N202  is  a  white  solid.  Its  solution  in  water  is  an  exceed- 
ingly feeble  acid.  The  warm  aqueous  solution  decomposes  slowly, 
giving  nitrous  oxide: 

H^NA-^HP  +  N2o, 

•% 

and  this  change  is  not  capable  of  reversal 

Nitrous  oxide  is  prepared  by  heating  ammonium  nitrate,  or  a  mix- 
ture of  a  salt  of  ammonium  and  a  nitrate: 

NH4NO3  ->  2H2O  +  N2O. 

The  steam  condenses,  and  the  nitrous  oxide  may  be  collected  over 
warm  water,  or  dried  and  compressed  into  steel  cylinders. 


OXIDES   AND   OXYGEN   ACIDS   OF    NITROGEN  301 

Its  solubility  in  cold  water  is  considerable:  at  0°,  130  volumes  in 
100;  at  25°,  60  in  100.  The  liquefied  gas  boils  at  -  89.8°  and  its 
vapor  tension  at  20°  is  49.4  atmospheres. 

A  glowing  splinter  of  wood  bursts  into  flame  when  introduced  into 
nitrous  oxide,  and  phosphorus,  sulphur,  and  other  combustibles  burn 
in  it  with  much  the  same  vigor  as  in  oxygen.  In  all  cases  oxides  are 
formed,  and  nitrogen  is  set  free. 

Metals  do  not  rust  in  nitrous  oxide,  and  the  haemoglobin  of  the 
blood  is  unable  to  use  it  as  a  source  of  oxygen.  It  is  employed  as  an 
anaesthetic  for  minor  operations.  The  symptoms  which  accompany 
its  administration  caused  it  to  receive  the  name  of  "  laughing  gas." 

Graphic  Formulce  of  Nitric  Acid  and  its  Derivatives  :  Ex- 
plosives. —  The  following  equation  for  the  formation  of  ammonium 
nitrate  by  neutralization  of  ammonium  hydroxide  with  nitric 
acid,  shows  the  graphic  (p.  156),  or  structural  formulae  of  these 
substances: 


H  H 

The  structural  formula  of  the  nitrate  is  intended  to  explain  the  fact 
that  the  salt  is  able  to  exist  at  all,  by  representing  the  oxygen  and 
hydrogen  as  being  separated  from  one  another  and  attached  to  differ- 
ent nitrogen  units.  When  the  equilibrium  of  the  system  is  disturbed 
by  heating,  the  oxygen  and  hydrogen  unite  to  form  water,  an  arrange- 
ment which  is  much  more  stable,  and  nitrous  oxide  (p.  300)  escapes 
with  the  steam. 

The  behavior  of  nitroglycerine  and  gun-cotton  (p.  294),  as  well  as 
of  ammonium  nitrite  (p.  279),  is  explained  in  the  same  way.  These 
substances  are  made  by  actions  which,  like  the  above  neutralization, 
take  place  in  the  cold,  and  the  groups,  containing  the  oxygen  on  the 
one  hand  and  carbon  and  hydrogen  on  the  other,  become  quietly 
united  without  more  serious  interaction.  When,  however,  the 
nitroglycerine,  for  example,  is  heated,  or  receives  a  mechanical 
shock,  the  carbon  and  hydrogen,  all  unite  with  the  oxygen  and  the 
nitrogen  escapes: 

4C3H6(N03)3  ->  12C02  +  10H20  +  6N2  +  02. 


302  COLLEGE   CHEMISTRY 

Active  ("Nascent")  Hitdrogen.  —  When  hydrogen  gas  is  led 
through  cold  nitric  acid,  little  or  no  action  occurs.  But  (p.  297) 
when  zinc,  or  some  other  metal  which  interacts  with  acids  to  give 
hydrogen,  is  placed  in  nitric  acid  the  latter  is  reduced.  To  explain 
the  apparent  greater  activity  of  the  hydrogen  in  the  second  instance, 
we  note  the  fact  that  it  is  liberated  on  the  surface  of  the  zinc.  The 
contact  effect  (catalytic  action)  of  the  zinc  increases  its  activity. 
Many  metals  have,  in  a  greater  or  less  degree,  this  power  of  increas- 
ing the  activity  of  hydrogen.  Thus,  hydrogen  absorbed  in  platinum 
or  palladium  (p.  73),  or  liberated  by  electrolysis  on  poles  made  of 
these  metals,  reduces  nitric  acid  readily.  Other  elements,  such  as 
the  chlorine  in  aqua  regia  (p.  299)  and  the  oxygen  in  making  sulphur 
trioxide  (p.  257),  are  also  rendered  more  active  by  contact  agents. 

This  more  active  state  of  hydrogen  is  described  as  the  nascent 
state,  because  it  happens  to  be  a  common  condition  of  hydrogen  when 
associated  with  substances  which  produce  it.  The  active  state  has, 
however,  no  necessary  connection  with  such  an  immediately  preced- 
ing act  of  liberation,  as  the  platinum  and  sulphur  trioxide  illustra- 
tions, and  the  following  experiment  show:  Three  test-tubes  are 
filled  with  very  dilute  potassium  permanganate  solution.  Zinc 
dust,  added  to  one,  generates  hydrogen  and  causes  decolorization. 
A  little  platinum  black  is  added  to  the  second,  and  hydrogen  gas 
is  led  through  this  and  the  third.  The  contact  action  of  the 
platinum  enables  the  hydrogen  quickly  to  induce  the  permanganate, 
while  the  third  portion  remains  unaltered. 

Exercises.  —  1.  Make  the  equation  for  the  interaction  of  ferrous 
chloride,  hydrochloric  acid,  and  nitric  acid  (p.  295),  and  for  all  the 
actions  concerned  when  the  test  for  a  nitrate  (p.  296)  is  applied  to 
sodium  nitrate.  What  volume  (at  0°  and  760  mm.)  of  NO  is 
obtained  from  one  formula- weight  of  nitric  acid  (p.  144)? 

2.  Make  correct  equations  for  the  formation  of  nitric  oxide  and 
nitrogen  tetroxide  by  the  action  of  carbon  on  nitric  acid  (p.  299). 

3.  Make  equations  for  the  interaction  of  iron  with  diluted  and 
with  concentrated  nitric  acid,  respectively  (p.  298).     The  iron  gives 
ferric  nitrate  Fe(N03)3. 

4.  Make  a  classified  list,  with  examples,  of  all  the  kinds  of  inter- 
actions which,  in  this  and  preceding  chapters,  have  been  named 
oxidations  and  reductions  (e.g.  pp.  52,  53,  75,  110,  112,  192,  212,  253, 
282,  297,  302.    See  also  Chemistry  of  copper  and  tin). 


CHAPTER  XXVII 
PHOSPHORUS 

Tlie  Chemical  Relations  of  the  Element.  —  There  are 
many  things  in  the  chemistry  of  phosphorus  and  its  compounds 
which  remind  us  of  nitrogen.  Yet  these  are  largely  referable  to  the 
fact  that  the  elements  are  both  non-metals  and  both  have  the  same 
valences,  viz.  three  and  five.  The  behavior  of  the  compounds  is 
often  very  different.  For  the  present  it  is  sufficient  to  say  that  both 
give  compounds  with  hydrogen,  NH3  and  PH3,  and  both  yield 
oxides  of  the  forms  X2O3,  X204,  and  X2O5.  The  first  and  last  of 
these  oxides  are  acid-forming,  and  phosphorus,  therefore,  gives  acids 
corresponding  to  nitrous  and  nitric  acids.  The  element  is  thus  a 
non-metal. 

Occurrence.  —  This  element  is  found  in  nature  in  the  form  of 
phosphates.  Calcium  phosphate  Ca3(PO4)2,  for  example,  occurs  in 
most  soils.  It  constitutes  a  large  part  of  the  solid  material  of  the 
bones  and  teeth  of  animals  and  of  the  beds  of  fossil  bones  found  in 
Florida  and  Tunis.  A  conspicuous  mineral  related  to  this  substance 
is  apatite,  Ca5F(PO4)3.  It  is  found  in  large  quantities  in  Canada, 
and  is  a  component  of  many  rocks.  Complex  organic  compounds 
containing  phosphorus  are  essential  constituents  of  protoplasm  and 
of  the  materials  of  the  nerves  and  the  brain. 

Preparation.  —  Brand,  merchant  and  alchemist,  of  Hamburg, 
discovered  phosphorus  (1669)  by  distilling  the  residue  from  evapo- 
rated urine,  in  the  course  of  his  search  for  the  philosopher's  stone. 
The  mode  of  preparing  it  from  bone-ash,  which  contains  83  per  cent 
of  calcium  phosphate,  was  first  published  by  Scheele  (1771).  Now 
the  less  expensive  calcium  phosphate  of  fossil  origin  is  employed. 

The  powdered  bone-ash  or  calcium  phosphate  and  sulphuric  acid 
(sp.  gr.  1.5)  are  heated  with  steam  and  stirred  in  a  wooden  vat: 
Ca3(P04)2  +  3H2SO4  ;±  ?H3P04  +  3CaS04J,. 

303 


304 


COLLEGE   CHEMISTRY 


The  calcium  sulphate  is  precipitated  during  the  heating,  and  sub- 
sequent concentration  of  the  filtrate.  This  syrupy,  crude  phosphoric 
acid  is  mixed  with  carbon  and  then  distilled  in  earthenware  retorts. 
Two  actions  take  place  in  succession.  The  phosphoric  acid  loses 
water  and  turns  into  metaphosphoric  acid,  then  the  latter  is  reduced 

by  the  carbon,  carbon  monoxide 
and  phosphorus  vapor  passing 
off: 


H3P04 
2HPO, 


30  +  HP03, 

+  6CO  +  2P. 


FIG.  57. 


A  pipe  from  the  tubular  clay 
retort  conducts  the  vapors  into 
cold  water,  in  which  the  phos- 
phorus collects. 

A  much  simpler  process  de- 
pends on  the  use  of  the  electric 
furnace  (Fig.  57).  The  calcium 
phosphate  is  mixed  with  the 
proper  proportions  of  carbon 
and  silicon  dioxide  (sand),  and 
the  mixture  is  introduced  con- 
tinuously into  the  furnace. 

The  discharge  of  an  alternating  current  between  carbon  poles  pro- 
duces the  very  high  temperature  which  the  action  requires.  The 
calcium  silicate  which  is  formed  fuses  to  a  slag,  and  can  be  with- 
drawn at  intervals.  The  gaseous  products  pass  off  through  a  pipe 
and  the  phosphorus  is  caught  under  water: 

,    Ca3(PO4)2  +  3SiO2  +  5C  -*  3CaSiO3  +  5CO  +  2P. 

We  may  regard  the  phosphate  as  being  composed  of  two  oxides, 
3CaO.P2O5.  It  thus  appears  that  the  calcium  oxide  has  united 
with  the  silica,  which  is  an  acid  anhydride  (cf.  p.  258) :  CaO  +  SiO2 
— >  CaSiO3,  while  the  phosphoric  anhydride  has  been  reduced. 

The  phosphorus,  after  purification,  is  cast  into  sticks  in  tubes  of 
tin  or  glass,  standing  in  cold  water. 

The  Electric  Furnace.  —  By  an  electric  furnace  is  understood 
an  electro-thermal  arrangement  in  which  the  heat  produced  by  some 
resistance  offered  to  the  current,  such  as  that  of  an  air-gap  between 


PHOSPHORUS  305 

the  carbons,  is  used  to  produce  chemical  change.  Electrolysis  plays 
no  part  in  the  phenomena,  and  an  alternating  current,  which  can 
produce  no  electrolytic  decomposition,  is  generally  employed.  The 
restricted  area  within  which  the  heat  is  developed  makes  possible 
the  attainment  of  a  high  temperature  (see  Calcium  carbide). 

Physical  Properties.  —  There  are  two  perfectly  distinct  kinds  of 
phosphorus,  known  as  ordinary,  or^ellpw  phosphorus,  and  red 
phosphorus.  Yellow  phosphorus,  prepared  as  described  above,  is  at 
first  transparent  and  colorless,  but  after  exposure  to  light  acquires  a 
superficial  coating  of  the  red  variety.  It  melts  at  44°  and  boils  at 
269°.  Its  molecular  weight  at  313°  is  128  and  the  formula,  therefore, 
P4.  Yellow  phosphorus  is  soluble  in  carbon  disulphide,  less  soluble 
in  ether,  and  insoluble  in  water.  It  is  exceedingly  poisonous,  less 
than  0.15  g.  being  a  fatal  dose.  Continued  exposure  to  its  vapor 
causes  necrosis,  a  disease  from  which  match-makers  are  liable  to 
suffer.  The  jawbones  and  teeth  are  particularly  liable  to  attack. 

Red  phosphorus  is  a  dull  red  powder  consisting  of  small  tabular 
crystals.  It  is  obtained  by  heating  yellow  phosphorus  to  about 
250°  in  a  vessel  from  which  air  is  excluded.  A  great  amount  of  heat 
is  evolved  in  the  transformation. 

Red  phosphorus  does  not  melt,  but  passes  directly  into  vapor.  Its 
vapor  is  identical  with  that  of  yellow  phosphorus.  It  is  insoluble  in 
carbon  bisulphide  and  other  solvents.  It  is  not  poisonous,  and, 
unlike  yellow  phosphorus,  does  not  require  to  be  kept  under  water 
to  avoid  spontaneous  combustion. 

Chemical  Properties.  —  Yellow  phosphorus  unites  directly  with 
the  halogens  with  great  vigor.  It  unites  slowly  with  oxygen  in  the 
cold,  and  with  sulphur  and  many  metals  when  the  materials  are 
heated  together.  The  slow  union  of  cold  phosphorus  with  atmo- 
spheric oxygen  is  accompanied  by  the  evolution  of  light,  although 
the  temperature  is  not  such  as  we  usually  associate  with  incan- 
descence. The  name  of  the  element  (Gk.  <££?,  light;  <#>epw,  bear) 
records  this  property.  Apparently  the  chemical  energy,  trans- 
formed in  connection  with  the  oxidation,  is  converted,  in  part  at 
least,  into  radiant  energy  instead  of  completely  into  heat.*  The 

*  The  same  production  of  light  from  chemical  action  in  a  cold  body  is  seen 
in  the  luminosity  of  certain  parts  of  some  animals,  such  as  fireflies  and  some 
species  of  fish.  In  many  violent  chemical  changes  the  light  given  out  is  con- 


306  COLLEGE   CHEMISTRY 

slow  oxidation  of  phosphorus  is  accompanied  by  the  production  of 
ozone,  but  the  nature  of  the  action  is  still  unknown  (cf.  p.  209). 

Red  phosphorus,  since  it  is  formed  with  evolution  of  heat,  contains 
less  energy  than  yellow  phosphorus  and  is  much  less  active.  It  does 
not  catch  fire  in  the  air  below  240°,  while  ordinary  phosphorus  ignites 
at  35-45°.  When  an  element  is  known  in  two  or  more  distinct 
forms,  as  is  the  case  with  oxygen,  phosphorus,  and  carbon  (q.v.), 
these  substances  are  called  allotropic  forms  of  the  element. 

Matches.  —  In  making  common  marches,  which  strike  on  any 
rough  surface,  the  splints  are  first  dipped  in  melted  sulphur  or  paraffin 
to  the  extent  of  about  half  an  inch.  The  head  is  often  composed  of 
manganese  dioxide  or  red  lead  and  a  little  potassium  chlorate,  which 
supply  oxygen,  a  small  proportion  of  free  phosphorus  (or  sulphide 
of  phosphorus)  and  antimony  trisulphide,  which  are  both  combus- 
tible, and  dextrin  or  glue. 

In  the  case  of  "  safety  "  matches,  the  mixture  upon  the  head  is  not 
easily  ignited  by  itself.  It  is  composed  of  potassium  chlorate  or 
dichromate,  some  sulphur  or  ^antimony  trisulphide,  and  a  little 
powdered  glass  to  increase  the  friction,  all  held  together  with  glue. 
Upon  the  rubbing  surface  on  the  box  is  a  thin  layer  of  antimony 
trisulphide  mixed  with  red  phosphorus  and  glue.  The  friction 
converts  a  little  of  the  red  phosphorus  into  vapor. 

Fhosphine.  —  Three  hydrides  of  phosphorus  are  known.  These 
are,  phosphine  PH3  (a  gas),  a  liquid  hydride  P2H4,  which  is  presum- 
ably the  analogue  of  hydrazine  (N2H4),  and  a  solid  hydride  P4H2. 

Phosphine  PH3  is  formed  slowly  by  the  action  of  active  hydrogen, 
from  zinc  and  hydrocjiloric  acid  at  70°,  upon  yellow  phosphorus. 
The  gas  may  be  made  by  boiling  yellow  phosphorus  with  potassium 
hydroxide  solution.  Potassium  hypophosphite  is  formed  at  the 
same  time:  -* 

3KOH  +  4P  +  3H20  -»  3KH2PO2  +PH3. 

The  gas  made  in  this  way  contains  a  little  of  the  vapor  of  the  liquid 

spicuously  greater  than  that  proper  to  the  temperature  produced  (c/.  p.  53), 
and  must  come,  therefore,  in  part,  directly  from  the  chemical  energy.  Thus, 
burning  magnesium  has  a  temperature  of  about  1350°,  while  the  production 
of  light  of  the  same  character,  by  mere  incandescence,  would  require  a  tem- 
perature of  about  5000°. 


PHOSPHORUS  307 

hydride,  which  is  spontaneously  inflammable,  and  consequently  the 
mixture  catches  fire  on  emerging  from  the  delivery  tube. 

The  simplest  method  of  preparing  phosphine  is  by  the  action  of 
water  upon  calcium  phosphide: 

Ca3P2  +  6H2O  -*  3Ca(OH)2  +  2PH3. 

This  action  is  analogous  to  that  of  water  upon  magnesium  nitride 
(p.  281),  by  which  ammonia  is  produced.  In  consequence  of  the 
fact  that  calcium  phosphide  is  a  substance  of..jfl%5ular  composition, 
a  mixture  of  all  three  hydrides  is  generally  oD^ained.  •  By  passing 
the  gas  through  a  strongly  cooled  delivery  tube,  however,  the  liquid 
compound  is  condensed  and  fairly  pure  phosphine  passes  on. 

Phosphine  is  a  colorless  gas,  whiqji  is  easily  decomposed  by  heat 
into  its  elements.  When  burned  it  forms  phosphoric  add.  It  is 
exceedingly  poisonous  and,  unlike  ammonia,  it  is  insoluble  in  water, 
and  produces  no  basic  compound  corresponding  to  ammonium 
hydroxide  when  brought  in  contact  with  this  substance.  It  resembles 
ammonia,  formally  at  least,  in  uniting  with  the  hydrogen  halides 
(see  below).  It  differs  from  ammonia,  however,  inasmuch  as  it  does 
not  unite  with  the  oxygen  acids.  Phosphine  acts  upon  solutions  of 
some  salts,  precipitating  phosphides  of  the  metals: 
3CuSO4  +  2PH3  ->  Cu3P2  +  3H2SO4. 

Phosphonium  Compounds.  —  Hydrogen  iodide  unites  with 
phosphine  to  form  a  colorless  solid  crystallizing  in  beautiful,  highly 
refracting,  square  prisms:  PH3  +  HI  — »  PH4L  Hydrogen  chloride 
combines  similarly  with  phosphine,  but  only  when  the  gases  are 
cooled  by  a  freezing  mixture,  or  are  brought  together  under  a  total 
pressure  of  18  atmospheres  at  14°.  When  the  pressure  is  released, 
rapid  dissociation  occurs. 

In  imitation  of  the  ammonia  nomenclature,  these  substances  are 
called  phosphonium  iodide  a»d  phosphonium  chloride  PH4C1.  They 
are  entirely  different,  however,  from  the  corresponding  ammonium 
derivatives,  for  the  PH4*  ion  is  unstable.  When  brought  in  contact 
with  water  they  decompose  into  their  constituents,  the  hydrogen 
halide  going  into  solution,  and  the  phosphine  being  liberated  as  a  gas. 

Halides  of  Phosphorus.  —  The  existence  of  the  following 
halides  has  been  proved  conclusively: 

....  P2I4  (solid) 

PF3  (gas)         PC13  (liquid)       PBr3  (liquid)        PI3   (solid) 
PF5  (gas)         PC15  (solid)         PBr5  (solid)          


308  COLLEGE  CHEMISTRY 

These  substances  may  all  be  formed  by  direct  union  of  the  elements. 
They  are  incomparably  more  stable  than  are  the  similar  compounds 
of  nitrogen.  They  are  all  hydrolyzed  by  water,  and  give  an  oxygen 
acid  of  phosphorus  and  the  hydrogen  halide  (see  below).  This 
action  was  used  in  the  preparation  of  hydrogen  bromide  (p.  162) 
and  hydrogen  iodide  (p.  167). 

Phosphorus  trichloride  PC13  is  made  by  passing  chlorine  gas  over 
melted  phosphorujkin  a  flask  until  the  proper  gain  in  weight  has 
occurred.  The  siWPance,  which  is  a  liquid  boiling  at  76°,  is  stable 
(cf.  p.  284).  When  excess  of  chlorine  is  employed,  phosphorus  pen- 
tachloride  PC15,  which  is  a  white  solid  body,  is  formed.  When  moist 
air  is  blown  over  any  of  these  substances,  the  water  is  condensed  to 
a  fog  by  the  hydrogen  halide.  In  the  case  of  the  interaction  of 
phosphorus  pentachloride  and  water,  phosphoric  acid  is  formed: 

PC15  +  4H2O  ->  H3P04  +  5HC1. 

Phosphorus  pentachloride,  when  heated,  reaches  a  vapor  tension 
of  760  mm.  at  140°,  and  while  still  solid.  It  therefore  passes  freely 
into  vapor  (boils,  so  to  speak)  at  this  temperature,  and  condenses 
directly  to  the  solid  form.  This  sort  of  distillation  is  called  sublima- 
tion. At  a  pressure  above  that  of  the  atmosphere  it  melts  at  148°. 
Partial  dissociation  occurs  in  the  vapor  (cf.  pp.  146,  181). 

Oxides  of  Phosphorus.  —  The  oxides  of  phosphorus  are  the 
so-called  trioxide  P4O6,  the  pentoxide  P2O5,  and  a  tetroxide  P2O4. 

The  pentoxide  is  a  white  powder  formed  when  phosphorus  is 
burned  with  a  free  supply  of  oxygen.  It  unites  with  water  with 
great  violence  to  form  metaphosphoric  acid  (see  below),  and  hence 
is  known  as  phosphoric  anhydride:  P205  +  H2O  — >  2HPO3.  In  the 
laboratory  this  action  is  frequently  utilized  for  drying  gases  (p.  287) 
and  for  removing  water  from  combination  (p.  294). 

The  trioxide  P4O6  is  obtained  by  burning  phosphorus  in  a  tube  with 
a  restricted  supply  of  air.  It  is  a  white  solid,  melting  at  22.5°  and 
boiling  at  173°.  This  oxide  is  the  anhydride  of  phosphorous  acid,  but 
it  unites  exceedingly  slowly  with  cold  water  to  form  this  substance, 
it  interacts  vigorously  with  hot  water,  but  phosphine,  red  phos- 
phorus, hypophosphoric  acid,  and  phosphoric  acid  are  amongst  the 
products,  and  very  little  phosphorous  acid  escapes  decomposition. 
When  this  oxide  is  heated  to  440°  it  decomposes,  giving  the  tetroxide 
P2O4  and  red  phosphorus. 


PHOSPHORUS  309 

Acids  of  Phosphorus.  —  There  are  six  different  acids  of  phos- 
phorus. Three  are  phosphoric  acids,  representing  the  same  stage 
of  oxidation  of  phosphorus,  but  different  degrees  of  hydration  of  the 
anhydride.  The  others  show  three  different  and  lower  states  of 
oxidation: 

Orthophosphoric  acid  H3PO4  (  =  3H2O,P2O5) 

Pyrophosphoric  acid  H4P2O7(  =  2H2O,P2O5) 

Metaphosphoric  acid  HPO3  (  =  H2O,P2O^) 

Hypophosphoric  acid  H4P2O6(  =  2H2O,P2O4) 

Phosphorous  acid  H3PO3  (  =  3H2O,P2O3) 

Hypophosphorous  acid  H3PO2  (  =  3H2O,P2O) 

T?ie  Phosphoric  Acids.  —  The  relation  between  the  three 
different  phosphoric  acids  may  be  seen  by  considering  them  as  being 
formed  from  phosphorus  pentoxide  (the  anhydride)  and  water.  In 
the  majority  of  cases  already  considered  this  sort  of  action  takes 
place  in  but  one  way.  Thus,  nitric  acid  is  known  in  but  one  form, 
which  is  produced  by  the  union  of  one  molecule  each  of  nitrogen 
pentoxide  and  water :  N2O5  +  H2O  — »  2HN03.  Similarly,  the  chief 
sulphuric  acid  is  the  one  formed  from  one  molecule  of  sulphur  trioxide 
and  one  molecule  of  water:  SO3  +  H2O  — >  H2SO4,  although  here  we 
have  also  disulphuric  acid  H2S2O7,  or  H2O,2S03. 

Now,  when  phosphoric  anhydride  acts  upon  water  we  obtain  a 
solution  which,  on  immediate  evaporation,  leaves  a  glassy  solid, 
HPO3,  known  as  metaphosphoric  acid.  This  is  H2O,  P205.  When, 
however,  the  solution  is  allowed  to  stand  for  some  days,  or  is  boiled 
with  a  little  dilute  nitric  acid,  whose  hydrogen-ion  acts  catalytically, 
the  residue  from  evaporation  is  H3PO4,  orthophosphoric  acid: 

P205  +  3H20  i-»  2H3PO4  or  HPO3  +  H2O  ->  H3PO4. 

This  acid  is  3H2O,  P2O5,  and  no  further  addition  of  water  can  be 
effected. 

Conversely,  when  orthophosphoric  acid  is  kept  at  about  255°  for  a 
time,  it  slowly  loses  water,  and  H4P2O7,  pyrophosphoric  acid,  is 

obtained:  2H,PO4 ->  H.O  +  H4P,Or 

This  acid  is  2H2O,  P205.  Further  desiccation  leaves  metaphos- 
phoric acid,  which  cannot  be  further  resolved  into  phosphorus 
pentoxide  and  water.  When  dissolved  in  water,  pyrophosphoric 


310 


COLLEGE   CHEMISTRY 


acid  slowly  resumes  the  water  which  it  has  lost  and  gives  the  ortho- 
acid  again.  The  pyro-acid  does  not  seem  to  be  formed  by  hydration 
of  the  meta-acid,  but  only  by  dehydration  of  the  ortho-acid.  The 
relations  of  all  these  substances  are  more  clearly  seen  in  the  graphic 
formulae: 


O 

0-H 
0  -H 
0-H 

0-H 
0-H 
O  -  H 
O 


0 

0-H 

O  -  H 

0 

O  -  H 

0-H 

O 


=  O 

-  O 

-  H 

f=0 

=  O 

P- 

-0 

<— 

0 

=  0 
-  0 

-  H 

P  ' 

[-0 

1=0 

=  0 

A  most  important  fact  to  be  noted  is  that  the  addition  or  removal  of 
water  leaves  the  valence  of  the  phosphorus  unchanged.  The  degree  of 
oxidation  of  the  phosphorus  and  its  valence  are  identical  in  the 
three  acids. 


The  Relations  of  Anhydrides  and  Oxygen  Acids.  —  Con- 
sidering the  anhydride  from  which  an  acid  is  derived  gives  us  the 
key  to  the  nature  and  behavior  of  the  acid.  It  tells  much  that 
the  formula  of  the  acid  does  not  tell.  For  example:  (1)  What  is  the 
valence  of  phosphorus  in  H3P03?  The  only  way  to  answer  this 
question  is  to  resolve  the  formula  (doubled,  if  necessary)  into  water 
and  the  anhydride,  3H2O,P2O3.  The  phosphorus  is  trivalent.  (2) 
How  is  this  acid  related  to  metaphosphoric  acid  HP03?  Resolve 
the  latter,  as  before,  H20,P2O5.  The  answer  is  that  in  the  latter 
the  phosphorus  is  quinquivalent.  (3)  How  can  we  get  phosphorous 
acid  from  metaphosphoric  acid,  or  vice  versa?  Considering  the 
anhydrides,  we  answer,  by  reduction  and  oxidation,  respectively. 
(4)  Is  pyrophosphoric  acid  H4P207,  because  it  contains  7O,  to  be 
made  from  all  others  by  oxidation?  Resolve  it  into  water  and 
anhydride,  2H20,P2O5.  We  then  perceive  that  to  make  it  from 
phosphorous  acid  requires  oxidation,  but  to  make  it  from  ortho-  or 
metaphosphoric  acid  requires  only  a  change  in  the  degree  of  hydra- 
tion: adding  or  subtracting  water,  since  it  adds  or  subtracts  hydrogen 
and  oxygen  in  equivalent  amounts,  is  not  oxidation  or  reduction. 
These  conceptions  have  been  used  before  (e.g.  pp.  260,  297,  309). 


PHOSPHORUS  311 

Considerations  like  the  foregoing  are  never  for  a  moment  absent 
from  the  mind  of  a  chemist  when  he  is  thinking  about  oxygen  acids. 
Every  such  acid  is  to  him  always  a  certain  anhydride,  plus  a  certain 
proportion  of  water  (see  Exercises  1  and  2).  He  applies  the  same 
method  to  salts,  subtracting  the  oxide  of  the  metal  (CaS04  = 
CaO,SO3)  to  determine  the  valence  and  state  of  oxidation  of  the 
non-metallic  element  in  the  compound. 

Orthophosphoric  Acid  JT3.PO4  and  Its  Salts. —  As  we  have 
seen  (p.  303) ,  ordinary  calcium  phosphate  is  the  source  of  the  impure, 
commercial  acid.  Pure  orthophosphoric  acid  may  be  made  by  boil- 
ing red  phosphorus  with  slightly  diluted  nitric  acid  and  evaporating 
the  water  and  excess  of  nitric  acid.  The  product  is  a  white,  crystal- 
line, deliquescent  solid. 

This  acid  is  much  weaker  than  sulphuric  acid,  and  is  dissociated 
chiefly  into  the  ions  H*  and  H2PO/.  The  dihydrophosphate-ion  is 
broken  up  to  some  extent  into  H*  and  HPO4",  as  we  learn  from  the 
fact  that  the  solution  of  the  sodium  salt  NaH2PO4  is  acid.  The  ion 
HP04"  is  hardly  dissociated  at  all,  for  a  solution  of  the  salt  Na^PC^ 
is  not  acid  in  reaction. 

As  a  tribasic  acid,  it  forms  salts  of  three  kinds,  such  as  NaH2P04, 
Na-jHPO^  and  Na3PO4.  These  are  known  respectively  as  primary, 
secondary,  and  tertiary  sodium  orthophosphate.  The  primary  sodium 
phosphate  is  faintly  acid  in  reaction.  The  secondary  one  is  slightly 
alkaline,  because  of  hydrolysis  arising  from  the  tendency  of  the 
hydrogen-ion  of  the  water  to  combine  with  theHPO4"to  formH2PO/, 
which  is  much  more  feebly  acid  than  is  phosphoric  acid  H3PO4. 
The  simplified  equation  (p.  254)  shows  the  reason  for  the  alkalinity 
of  the  solution:  HPO4"  +  H'  +  OH/-^H2PO/  +  OH'.  The  ter- 
tiary phosphate  is  stable  only  in  solid  form,  and  can  be  made  by 
evaporating  to  dryness  a  mixture  of  the  secondary  phosphate  and 
sodium  hydroxide: 

Na2HP04  +  NaOH  <=>  Na3PO4  +  H2O  |. 

When  the  product  is  dissolved  in  water,  the  action  is  reversed  (cf. 
p.  253).  Mixed  phosphates  are  also  known,  particularly  sodium-am- 
monium phosphate  (microcosmic  salt)  NaNH4HP04,  4H2O,  and  the 
insoluble  magnesium-ammonium  phosphate  MgNH4PO4.  Primary 
calcium  phosphate  (q.v.),  known  in  commerce  as  "  superphosphate," 
is  used  as  a  fertilizer. 


312  COLLEGE   CHEMISTRY 

The  tertiary  phosphates  are  unchanged  by  heating.  The  primary 
and  secondary  phosphates,  however,  retaining,  as  they  do,  some  of 
the  original  hydrogen  of  the  phosphoric  acid,  are  capable  of  losing 
water  like  phosphoric  acid  itself,  when  heated.  The  actions  are 
slowly  reversed  when  the  products  are  dissolved  in  water: 

NaH2P04  <±  NaP03    +  H20, 
=±  Na4P2O7  +  H2O. 


It  will  be  seen  that  the  meta-  and  pyrophosphates  of  sodium  are 
formed  by  these  actidns;  and  this  is  indeed  the  simplest  way  of  form- 
ing these  substances,  since  the  acids  themselves  are  not  permanent  in 
solution,  and  are  too  feeble  to  lend  themselves  to  exact  neutraliza- 
tion. Ammonium  salts  of  phosphoric  acid  lose  ammonia,  as  well 
as  water,  when  heated  (cf.  p.  283).  Thus,  microcosmic  salt  gives 
primary  sodium  phosphate: 

NaNH4HP04  -»  NH3  +  NaH2PO4  -»  NaPO3  +  H2O, 
and  this  in  turn  is  converted  into  the  metaphosphate  by  loss  of  water. 

Pyrophosphoric  Acid  and  Metaphosphoric  Acid.  —  Pyro- 
phosphoric  acid  H4P2O7,  although  tetrabasic,  gives  only  the  neutral 
salts,  such  as  Na4P207,  and  those  in  which  one-half  of  the  hydrogen 
has  been  displaced  by  a  metal,  such  as  Na2H2P207. 

Metaphosphoric  acid  HPO3  is  the  "  glacial  phosphoric  acid  "  of 
commerce,  and  is  usually  sold  in  the  form  of  transparent  sticks.  It 
is  obtained  by  heating  orthophosphoric  acid,  or  by  direct  union  of 
phosphorus  pentoxide  with  a  small  amount  of  cold  water.  It  passes 
into  vapor  at  a  high  temperature,  and  its  vapor  density  corresponds 
to  the  formula  (HPO3)2. 

Sodium  metaphosphate  NaP03,  in  the  form  of  a  small  globule 
obtained  by  heating  microcosmic  salt  on  a  platinum  wire,  is  used  in 
analysis.  When  minute  traces  of  oxides  of  certain  metals  are  placed 
upon  such  a  globule,  known  as  a  bead,  and  heated  in  the  Bunsen 
flame,  the  mass  is  colored  in  various  tints  according  to  the  oxide  used 
(bead  test)  .  This  action  may  be  understood  when  we  consider  that 
sodium  metaphosphate  takes  up  water  to  form  primary  sodium 
orthophosphate:  NaPO3  +  H2O—  >  NaH2PO4.  In  the  same  way, 
but  at  higher  temperatures,  it  is  able  to  take  up  oxides  of  elements 
other  than  hydrogen,  giving  mixed  orthophosphates.  Thus,  with 


PHOSPHORUS 


oxide  of  cobalt  a  part  of  the  metaphosphate  unites  according  to 
equation: 


and  the  product  confers  a  blue  color  on  the  bead. 

Distinguishing  Tests.  —  When  a  solution  of  nitrate  of  silver  is 
added  to  a  solution  of  orthophosphoric  acid  or  any  soluble  orthophos- 
phate,  a  yellow  precipitate  of  silver  orthophosphate  Ag3PO4  is  pro- 
duced. This  is  a  test  for  phosphate-ion.  With  pyrophosphoric 
acid  or  any  pyrophosphate  the  product  is  white  Ag4P207.  With 
metaphosphoric  acid  a  white  precipitate,  AgPO3,  is  obtained  also. 
Metaphosphoric  acid  coagulates  a  clear  solution  of  albumen,  while 
pyrophosphoric  acid  has  no  visible  effect  upon  it. 

Phosphorous  Acid  JT3.PO3.  —  When  added  to  cold  water  phos- 
phorus trioxide  (P4O6)  yields  phosphorous  acid  very  slowly.  With 
hot  water  the  action  is  exceedingly  violent  and  complex  (p.  308). 
This  acid  may  be  obtained  also  by  the  action  of  water  upon  phos- 
phorus trichloride,  tribromide  (p.  162),  or  tri-iodide  and  evaporation 
of  the  solution: 


A  certain  amount  of  this  acid,  along  with  phosphoric  acid  and  hypo- 
phosphoric  acid,  is  formed  when  moist  phosphorus  oxidizes  in  the 
air. 

In  spite  of  the  presence  of  three  hydrogen  atoms,  this  acid  is 
dibasic,  and  two  only  are  replaceable  by  metals.  To  express  this 
fact,  the  first  of  the  following  formula  is  preferred: 


since  the  symmetrical  formula  would  indicate  no  difference  between 
the  three  hydrogen  atoms.  H  united  directly  to  P,  as  here  and  in 
PH3,  is  not  acidic.  Phosphorous  acid  is  a  powerful  reducing  agent, 
precipitating  silver,  for  example,  in  the  metallic  form  from  solutions 
of  its  salts.  When  heated,  it  decomposes,  giving  the  most  stable 
acid  of  phosphorus  (cf.  pp.  191,  196,  199,  266,  299),  namely,  meta- 
phosphoric acid,  and  phosphine: 

4H3PO3  -»  3HP03  +  3H2O  +  PH3. 


314  COLLEGE   CHEMIS1RY 

Sulphides  of  Phosphorus.  —  Yellow  phosphorus  when  heated 
with  sulphur  unites  with  explosive  violence.  By  using  red  phos- 
phorus the  action  can  be  controlled.  By  employing  the  proper 
proportions  the  pentasulphide  P2S5  is  secured.  It  is  purified  by  dis- 
tillation, and  is  a  yellow  crystalline  mass  (m.-p.  274°,  b.-p.  530°). 
Phosphorus  pentasulphide  is  hydrolyzed  by  cold  water: 
P2S5  +  8H2O  ->  2H3PO4  +  5H2S. 

Other  sulphides,  ^483,  P2S3,  and  P3S6,  may  be  prepared  by  using 
the  constituents  in  the  proportions  represented  by  these  formulae. 

Comparison  of  Phosphorus  with  Nitrogen  and  with  Sul- 
phur. —  Although  phosphorus  and  nitrogen  are  regarded  as  belong- 
ing to  one  family,  the  differences  between  them  are  more  conspicuous 
than  the  resemblances.  The  latter  are  confined  almost  wholly  to 
matters  concerned  with  valence.  The  differences  are  seen  in  the 
facts  that  nitrogen  is  a  gas  and  exists  in  but  one  form,  while  phos- 
phorus is  a  solid  occurring  in  two  varieties,  and  that  the  former  is 
inactive  and  the  latter  active.  The  contrasts  between  phosphine 
and  ammonia  (p.  307)  and  between  the  halides  of  the  two  elements 
(pp.  284,  308)  have  been  noted  already.  The  pentoxide  of  nitrogen 
decomposes  spontaneously;  that  of  phosphorus  is  one  of  the  most 
stable  of  compounds.  Nitric  acid  is  very  active;  the  phosphoric 
acids  are  quite  the  reverse. 

On  the  other  hand,  the  resemblance  of  phosphorus  to  sulphur  is 
marked.  Both  are  solids,  existing  in  several  forms.  Both  yield 
stable  compounds  with  oxygen  and  chlorine.  The  hydrogen  com- 
pounds interact  with  salts  to  give  phosphides  of  metals  and  sulphides 
of  metals,  respectively.  Against  these  must  be  set  the  facts,  that 
hydrogen  sulphide  does  not  unite  with  the  hydrogen  halides  at  all 
while  phosphine  gives  the  phosphonium  halides,  and  that  phosphoric 
acid  is  hard  to  reduce  while  sulphuric  acid  is  reduced  with  compara- 
tive ease. 

Eocercises. —  1.  What  are  the  valences  of  the  non-metals  in: 
H2S207,  H2Cr207,  KMn04,  KH2PO2,  H3N04,  NaH2PO3,  Na2P2Oe. 
Name  these  substances. 

2.  Is  it  oxidation  or  reduction,  or  neither,  when  we  make,  (a) 
N3O4  from  HNO3,  (b)  S02  from  H2SO3,  (c)  HPO3  from  H3P03,  (d) 
H2S207  from  H2SO4,  (e)  Na,SO4  from  NaHSO3? 


PHOSPHORUS  315 

3.  Why  would  a  mixture  of  potassium  dichromate  and  hydro- 
chloric acid  (p.  253)  be  less  suitable  than  nitric  acid,  as  an  oxidizing 
agent  for  making  phosphoric  acid  from  red  phosphorus? 

4.  Why  is  not  the  tertiary  phosphate  of  sodium  (p.  311)  decom- 
posed by  heating?    What  tertiary  phosphates  would  be  decomposed 
by  this  means? 

5.  Formulate  the  hydrolyses  of  the  secondary  and  tertiary  sodium 
orthophosphates  as  was  done  for  sodium  sulphide  (p.  253). 

6.  How  should  you  prepare  Cs^P2O7  and  Ca(PO3)2? 

7.  What  product  should  you  confidently  expect  to  find  after 
heating  (p.  313),  (a)  sodium  phosphite,  Na^HPOg,  (b)  potassium 
hypophosphite?    Make  the  equations. 

8.  Compare  the  elements   chlorine   and  phosphorus   after  the 
manner  of  the  comparisons  on  p.  314. 


CHAPTER  XXVIII 
\t  CARBON  AND  THE  OXIDES  OF  CARBON 

,  The  Chemical  Relations  of  the  Element.  —  The  elements  of 
the  carbon  family  (p.  276)  are  carbon,  silicon,  germanium,  tin,  and 
lead.  Of  these  the  first  two  are  entirely  non-metallic,  while  the 
others  are  metallic  elements  showing  more  or  less  strong  resemblances 
to  the  non-metals.  All  these  elements  are  quadrivalent.  With  the 
exception  of  silicon,  however,  they  all  form  compounds  in  which 
they  are  bivalent. 

On  account  of  the  fact  that  the  majority  of  the  substances  com- 
posing, and  produced  by,  living  organisms  are  compounds  of  carbon, 
the  chemistry  of  these  compounds  is  known  as  organic  chemistry. 
It  was  at  first  supposed  that  the  artificial  production  of  such  com- 
pounds, e.g.  without  the  intervention  of  life,  was  impossible.  But 
many  natural  organic  products  have  now  been  made  from  simpler 
ones  or  from  the  elements,  a  process  called  synthesis,  and  the  prepa- 
ration of  the  others  is  delayed  only  in  consequence  of  difficulties 
caused  by  their  instability  and  complexity.  On  the  other  hand, 
hundreds  of  compounds  unknown  to  animal  or  vegetable  life,  includ- 
ing many  valuable  drugs  and  dyes,  have  now  been  added  to  the 
catalogue  of  chemical  compounds. 

The  elements  entering  into  carbon  cooij$unds  are  chiefly  hydrogen 
and  oxygen.  After  these,  nitrogen,  phosphorus,  the  halogens,  and 
sulphur  may  be  named. 

CARBON  C. 

Occurrence.  —  Large  quantities  of  carbon  are  found  in  the  free 
condition  in  nature.  The  diamond  is  the  purest  natural  carbon. 
Graphite,  or  plumbago,  which  is  the  next  purest,  is  found  in  limited 
amounts,  and  is  a  valuable  mineral.  Coal  occurs  in  numerous  forms 
containing  greatly  varying  proportions  of  free  carbon.  Small  quan- 
tities of  the  free  element  have  been  found  in  meteorites. 

In  combination,  carbon  is  found  in  marsh-gas,  or  methane  CH4, 

816 


CARBON   AND  THE   OXIDES  OF   CARBON 


317 


FIG.  58. 


which  is  the  chief  component  of  natural  gas.  The  mineral  oils  con- 
sist almost  entirely  of  mixtures  of  various  compounds  of  carbon  and 
hydrogen.  Whole  geological  formations  are  composed  of  carbonates 
of  common  metalif,  particularly  calcium  carbonate  or  limestone, 
and  a  double  carbonate  of  calcium  and  magnesium,  known  as 
dolomite. 

The  Diamond.  —  The  allotropic  forms  (p.  306)  of  c^Bbn  differ 
very  markedly  in  their  physical  properties.     Diamonds,  which  are 
found  in  India,  Borneo,  Brazil,  and  South  Africa,  are  scattered 
sparsely  through  metamorphic  and 
volcanic  rocks  which  seem  to  have 
undergone       secondary      changes. 
They   are   covered    with    a    crust 
which  entirely  obscures  their  luste^, 
and    possess     natural     crystalline 
forms    belonging    to    the    regular 
system.     A   form   related    to    the 
octahedron    (p.  94)   is  frequently 

observed.  It  should  be^  noted  that  this  natural  form  bears  no 
relation  whatever  to  me  pseudo-crystalline  shape  which  is 
conferred  upon  the  stone  by  the  diamond-cutter.  Thus,  a 
"  brilliant "  possesses  one  rather  large,  flat  face,  which  forms  the 
base  of  a  many  sided  pyramid  (Fig.  58,  showing  two  views).  This 
form  is  given  to  the  stone,  in  order  that  the  maximum  reflection 
of  light  from  its  interior  may  be  produced.  The  diamond  is  harder 
than  any  other  variety  o^natter,  with  the  exception,  perhaps,  of 
one  compound  of  boron,  «hBk  only  one  or  two  other  materials,  like 
carborundum,  approach^.  Rs  specific  gravity  is  about  3.5,  and  it 
is  the  densest  form  of  carbon:  Few  materials  are  capable  of  dissolv- 
ing any  of  the  forms  of  carbon.  Molten  iron  (q.v.)  dissolves  five  or 
six  per  cent,  part  of  which  goes  into  combination;  but  usually  only 
graphite  is  found  in  the  cooled  product.  Moissan  (1887),  however, 
succeeded  in  preparing  microscopic  fragments  of  diamonds  in  this 
way.  The  diamond  is  a  nonconductor  of  electricity. 

The  largest  diamond  known,  the  Cullinan,  weighs  621  g.  (uncut). 
The  Kohinoor  weighs  22  g. 

That  the  diamond  contains  nothing  but  carbon,  is  shown  by  the 
fact  that  when  burned  it  produces  nothing  but  carbon  dioxide. 


318 


COLLEGE   CHEMISTRY 


Graphite. — Graphite  (Gk.  ypa<#>o>;  I  write)  is  found  in  Cumberland, 
Siberia,  and  Ceylon.  The  form  appears  to  belong  to  the  hexagonal 
system.  The  mineral,  is  extremely  soft,  In  utter  cpntrast  to  the 
diamond,  and  has.a  smaller  .specific  gravity  (abo€t  2.3).  It  conducts 
electricity.  It  is  made .  artificially  by  an  electro-thermal  <  process 
(cf.  p.  304).  A  powerfuj  alternating  current  is  passed  through  a 
mass  of  annular  anthracite,  and  the  latter,  although  not.riielted  by 
the  higl^Pmperature/is  largely  converted  into  graphite. 

Graphite  is  ftw  used,  exclusively  for  making  the  anodes  in  the 

•lectrolytic  manufacture  of  chlorine  and  in  related  processes.     lyttxed 

with  fine  clay  it  forms  the  "  lead  "  of  lead  pencils.   As  "  black-lead  ", 

ft  is  employed  to  protect  ironware  fipm  rusting.     It  is  used  as  a 

lubricant  in  cases  where  oil  would  be  decomposed  by  heat. 

Amorphous  Carbon.  —  This  «s  the  name  given  to  the  varieties  of 
the  element  which  have  no  Crystalline  form.  They  vary  in  specific 
gravity  up  to  1^9.  Coke,  which  is  now  manufactured  in  immense 
quantities  by  heating  coal  until  all  the  volatile  matter  has  been  dis- 
tilled off,  is  a  very  dense  variety  of  amorphous  carbon*  It  is  used  in 
the  reduction  of  iron  ores.  f 

By  the  imperfect  combustion  of  heavy  oils  and  resins,  in  which 
the  flame  plays  upon  a  cooled  surface,  a  finely  divided  form  of  carbon, 
lampblack,  is  produced.  This  is  used  in  making  printer's  ink. 

Charcoal  is  chiefly  made  by  the  heating  of  wood  out  of  contact  with 
air.  In  the  more  refined  forms  of  the  process  the  charring  is  con- 
ducted in  retorts,  and  the  materials  which  distil  off  are  used  in  vari- 
ous' ways.  Wood  consists  largely  of  c^Wose  (C6H1005)n,  incrusted 
with  lignin  and  holding  much  moistu™  .••  resinous  material.  The 
products  of  its  distillation  are  partly  J|  PUS  and  partly  fluid.  The 
gases,  consisting  mainly  of  hydrogen,  methane  CH4,  ethane  C2H6, 
ethylene  C2H4,  and  carbon  monoxide  CO,  are  employed,  on  account 
of  their  combustibility,  as  fuel  in  the  distillation  itself.  The  fluids 
form  a  complex  mixture  containing  large  quantities  of  water,  methyl 
alcohol  CH3OH  (wood  spirit),  acetic  acid,  acetone  (CH3)2CO,  and  tar. 
Wood  charcoal  exhibits  the  cellular  structure  of  the  material  from 
which  it  was  made,  and  is  therefore  highly  porous.  When  the 
charcoal  is  burned,  the  mineral  constituents  of  the  wood  appear  in 
the  ash.  ^ 

For  certain  purposes,  charcoMs  made,  in  the  same  fashion  as  the 


• 

CARBON   AND  THE   OXIDES   OF   CARBON  319 

» 

above,  froq|  bones  and  from  blood,  find  wide  application.  The 
former,  callecf  bone  black,  contains  much  calcium  phosphate  (p.  803). 
In  the  chemical  laboratory,  pure  carbon  is  made,  as  a  rule,  by  the 
charring  of  sugar  (cane-sugar,  C^H^Ou). 

The  tendency  of  almost  all  carbon  compounds  to  char,  when 
heated,  is  used  as  a  means  of  recognizing  their  presence. 

Properties  of  Charcoal.  —  Charcoal- exhibits  cert  a i Properties 
which  are  not  shared  to  any  extent  by  other  format  can5on.  For 
example,  it  can  take  up  large  quantities  of  many  gases.  BoxwocA 
charcoal  will  in  this  way  absorb  ninety  times  its  own  volume  or 
ammonia,  fifty-five  yplumes  ,°£]jjj|^)gen  sulphide,  or  nine  volumes  of 
oxygen.  Freshly,  made  oogwoocF  charcoal  (used  in  making  the 
best  gunpowder),  when  pulverized  immediately  after  its  preparation, 
often  catches  fire  spontaneously  on  account  of  the  heat  liberated  by 
the  condensation  of  oxygen.  It  is  therefore  set  aside  for  two  weeks, 
to  permit  the  slow  absorption  of  mofstufe  and  air,  before  being 
ground  up.  The  absorbed  gases  may  be  removed  unchanged  by 
\  heating  the  charcoal  in  a  vacuum.  The  disappearance  of  these 
immense  quantities  of  gas  into  small  pieces  of  charcoal  is  described 
as  adsorption,  and  is  caused  by  the  cohesion  of  the  gases  to  the  very 
extensive  internal  surface  which  the  charcoal  possesses.  Glass  shows 
the  same  property,  though  in  a  smaller  degree  (p.  80).  Solid  and 
liquid  bodies  are  also  ia  many  cases  l!aken  up  by  charcoal  in  a  similar 
fashion.  Thus,  strychnine  may  be  removed  from  an  aqueous  solu- 
tion by  agitation  of  the  latter  with  charcoal.  In  the  manufacture  of 
whiskey  (q*v.),  the  fusel  oil,  which  is  extremely  harmful,  is  in  many 
cases  removed  by  filtratio^of  the  diluted  spirit  through  charcoal, 
before  rectification.  Orgamc  coloring  matters,  such  as  litmus  and 
indigo,  belong  to  the  class  of  bodies  thus  extracted  from  solution  by 
charcoal.  In  the  refining  of  sugar  the  syrup  is  boiled  with  charcoal 
for  the  purpose  of  removing  a  brown  resin,  in  order  that  the  product 
may  be  perfectly  white.  It  is,  in  part,  upon  tiris  property  that  we 
rely,  also,  in  the  employment  of  charcoal  nrfeers.  The  organic 
materials  dissolved  in  the  drinking  water  undergo  adsorption  in  the 
charcoal.  In  this  connection,  however,  it  must  be  remembered  that 
the  quantity  which  a  given  mass  of  charcoal  may  take  up  is  limited, 
and  that  careful  cleansing  is  requiij^  in  order  that  the  efficiency  of 
the  filter  may  be  maintained. 


[uir|i  : 


320 


COLLEGE    CHEMISTRY 


Coal.  —  Peat,  brown  coal,  soft  coal,  and  anthracite  represent,  in 
a  general  way,  ^different  stages  in  the.  decomposition*  01  vegetable 
matter  in  absence  of  air.  Water  and  compounds  of  carbon  and 
hydrogen  are  given  off  in  the  process.  The  following  table  shows 
this  change  in  composition  and  the  relations  of  the  substances  to 
fresh  wood  on  the  one  hand  and  charcoal  and  coke  on  the  other: 


—  ~~A  — 

Percentage,  excluding  Ash 
and  Moisture,  of  : 

Percentage 
Ash. 

Percentage 
of  Water 
(Air  Dried). 

Calorific 
Value 
per  g.  in 
Calories. 

« 

H 

O 

-ii 

48 
32 
24 
12 

5.4 

2.5 

N 

2 
1 
1 

1 

Wood    .... 
Peat 

45 
60 
70 
82 
94 
95 
96 

6 
6 
5 
5 
3 
1.7 
0.7 

1.5 
5-20 
3-30 
1-15 
1.5 
4 
3.4-11 

18-20 
20-30 
15 
4 
2 
6.5 
2 

2700 
3500 
30-6000 
66-8000 
70-8000 
8080 
7700 

Brown  coal     . 
Soft  coal  .    .    . 
Anthracite  .    . 
Charcoal  .    .    . 
Coke     . 

Chemical  Properties  of  Carbon.  —  The  most  common  uses 
of  carbon  depend  upon  its  great  tendency  to  unite  with  oxygen, 
forming  carbon  dioxide.  Under  some  circumstances  carbon  mon- 
oxide is  produced.  Aside  from  the  direct  employment  of  this  action 
for  the  sake  of  the  heat  which  is  liberated,  it  is  used  also  in  the 
reduction  of  ores  of  iron,  copper,  zinc,  and  many  other  metals. 
When,  for  example,  finely  powdered  cupric  oxide  and  carbon  are 
heated,  copper  is  obtained.  The  gas  given  off  is  either  carbon 
dioxide,  or  a  mixture  of  this  with  carjj^m  monoxide,  according  to 
the  proportion  of  carbon  used: 

CuO  +  C  -»  Cu  +  CO, 
2CuO  +  C  ->  2Cu  +  CO2. 

Carbon  unites  directly  with  hydrogen  very  reluctantly.  When  an 
electric  arc  is  produced  between  carbon  poles  in  a  tube  through 
which  hydrogen  passes,  traces  of  acetylene  C2H2  are  formed. 

At  the  high  temperatures  produced  in  the  electric  furnace,  carbon 
unites  with  many  metals  and  some  non-metals.  Compounds  formed 
in  this  way  are  known  as  carbicfe  such  as  aluminium  carbide  A14C3, 
calcium  carbide  CaC2,  and  caiMrundum  CSi. 


CARBON  AND  THE   OXIDES   OF   CARBON 


321 


Calcium  Carbide  CaC2.  —  This  compound,  which  is  colorless 
when  pure,  is  manufactured  in  an  electric  furnace,  by  the  interaction 
of  finely  pulverized  limestone  or  quicklime  with  coke: 
CaO  +  3C  -»  CaC2  +  CO. 

The  operation  is  a  continuous  one,  the  materials  being  thrown  into  the 
left  side  of  the  drum  (Fig.  59,  diagrammatic),  and  the  product 
removed  on  the  right.  The  carbon  poles  are  fixed.  The  dfe  having 
been  established,  the  drum 
is  rotated  slowly  as  the  car- 
bide accumulates.  The  cur- 
rent enters  by  one  carbon, 
passes  through  the  carbide, 
and  leaves  by  the  other. 
The  high  resistance  of  the 
partially  transformed  ma- 
terial causes  the  production 
of  the  heat.  When  the  action 
in  one  layer  approaches 
completion,  the  resistance 
falls,  the  current  increases, 
and  an  armature  round 
which  the  wire  passes  (not 
shown  in  Fig.  59)  comes  into 
operation  and  turns  the 

drum.  In  this  way  the  carbide  just  formed  is  continuously  moved 
away  from  the  carbons,  and  new  material,  introduced  on  the  left, 
falls  into  the  path  of  the  current.  The  iron  plates  which  form  the 
circumference  of  the  drum  are  added  on  the  left  and  removed  on  the 
right,  where  also  the  carbide  is  broken  out  with  a  chisel.  The  drum 
revolves  once  in  about  three  days.  The  product  is  used  for  making 
acetylene  (q.v.). 

CARBON  DIOXIDE  AND  CARBONIC  ACID. 

Occurrence.  —  Carbon  dioxide  is  present  in  the  atmosphere,  and 
issues  from  the  ground  in  large  quantities  in  certain  neighborhoods, 
as,  for  example,  near  the  Lake  of  Laach,  in  the  so-called  Valley  of 
Death  in  Java,  and  in  the  Grotta  del* Cane  near  Naples.  Efferves- 
cent mineral  waters  contain  it  in  solution,  and  their  effervescence  is 


FIG.  59. 


822  COLLEGE   CHEMISTRY 

caused  by  the  escape  of  the  gas  when  the  pressure  is  reduced.  Well- 
known  waters  of  this  kind  are  those  of  Selters  and  of  the  Geyser 
Spring  at  Saratoga. 

Modes  of  Formation.  —  Carbon  dioxide  is  produced  by  combus- 
tion of  carbon  in  the  presence  of  an  excess  of  oxygen : 

%  C  +  O2  -*  CO2. 

The  combustion  of  all  compounds  of  carbon,  as  well  as  the  slow  oxida- 
tion  in  the  tissues  of  plants  and  animals,  yield  the  same  product. 

It  was  Joseph  Black  (1757)  who  first  recognized  the  gas  as  a  dis- 
tinct substance.  He  observed  its  formation  when  marble  or  mag- 
nesium carbonate  was  heated: 

CaCO3  <=>  CaO  +  CO2, 

and  named  the  gas  "  fixed  air  "  from  the  fact  that  it  was  contained 
in  these  solids.  The  above  action  had  been  used  for  centuries  in 
making  quicklime  (calcium  oxide).  All  common  carbonates,  except- 
ing the  normal  carbonates  of  potassium  and  sodium,  decompose  in 
this  way,  leaving  the  oxide  of  the  metal. 

Black  found  that  the  gas  was  also  produced  when  acids  acted  upon 
carbonates,  and  this  method  is  commonly  employed  in  the  laboratory : 

CaCO3  (solid)  £3  CaCO3  (dslvd)  <->  Ca"  +  CO3"\  «_„  rn  ^_  „  n       pn 
2HC1  (dslvd)  £=>  2C1'  +  2H'    f  *=*H9C°t  _>  H2O        ,O2. 

Since  the  carbonic  acid  is  very  slightly  ionized,  the  action  is  like  that 
of  acids  on  sulphites  (p.  256).  The  carbonate  of  calcium,  however,  is 
very  slightly  soluble,  so  that  an  additional  equilibrium  controls  its 
solution.  In  this  respect  the  action  is  like  that  of  acids  on  ferrous 
sulphide  (p.  251). 

Carbon  dioxide  is  formed  in  decay  (p.  287)  and  in  fermenta- 
tion (q.v.). 

Physical  Properties.  —  Carbon  dioxide  is  a  colorless,  odorless 
gas.  It  is  heavier  than  air.  The  G.M.V.  weighs  44.26  g.  Liquefied 
carbon  dioxide  boils  at  —  79°.  The  sp.  gr.  of  the  liquid  at  0°  is  0.95. 
At  0°  its  vapor  tension  is  35.4  atmospheres  and  at  20°  59  atmos- 
pheres. It  must  be  preserved,  therefore,  in  very  strong  cylinders 
of  mild  steel.  Large  quantities  of  it,  often  collected  from  fermen- 
tation vats,  are  sold  in  such  cylinders,  and  used  in  operating  beer- 
pumps  and  in  making  aerated  waters.  When  the  liquid  is  allowed 


CARBON   AND   THE   OXIDES   OF   CARBON  323 

to  flow  out  into  an  open  vessel  it  cools  itself  by  its  own  evaporation 
and  forms  a  white,  snowlike  mass.  Solid  carbon  dioxide  evaporates 
without  melting  (cf.  p.  308). 

Carbon  dioxide  gas  (760  mm.  and  15°)  dissolves  in  its  own  volume 
of  water.  Up  to  four  or  five  atmospheres  Henry's  law  describes  its 
solubility  accurately.  An  aqueous  solution,  under  a  pressure  of  2-3 
atmospheres,  is  familiarly  known  as  soda-water. 

Chemical  Properties.  —  Carbon  dioxide  is  a  stable  compound. 
At  2000°,  the  dissociation  reaches  7.5  per  cent,  or  about  four  times 
that  of  water:  2C02  <=±  2CO  +  02. 

The  more  active  metals,  like  magnesium,  burn  brilliantly  when 
ignited  in  a  cupel-shaped  lump  of  solid  carbon  dioxide,  producing 
the  oxide  and  free  carbon.  Less  active  metals,  such  as  zinc  and  iron, 
give  an  oxide  and  carbon  monoxide  (q.v.). 

Carbon  dioxide  unites  directly  with  many  oxides,  particularly  those 
of  the  more  active  metals,  such  as  the  oxides  of  potassium,  sodium, 
calcium,  etc.  Hence  the  decomposition  of  calcium  carbonate  by 
heating  (p.  322)  is  a  reversible  action. 

Carbon  dioxide,  when  dissolved  in  water,  forms  an  unstable  acid: 
H20  -f  CO2  <=±  H2CO3. 

The  name  carbonic  acid  is  frequently,  though  improperly,  given  to 
the  anhydride  CO2,  which  has  no  acid  properties. 

Chemical  Properties  of  Carbonic  Acid  Jf2CO3. — The  solu- 
tion of  carbon  dioxide  in  water  exhibits  the  properties  of  a  weak  acid. 
It  conducts  electricity,  although  not  well.  It  turns  litmus  red. 
The  ionization  takes  place  chiefly  according  to  the  equation: 

H2CO3  <=±  H'  +  HCCy. 

Carbonates  and  Bicarbonates  —  When  excess  of  an  aqueous 
solution  of  carbonic  acid  is  mixed  with  a  solution  of  a  base  like  sodium 
hydroxide,  or,  as  the  operation  is  more  usually  performed,  when 
carbon  dioxide  is  passed  into  a  solution  of  the  alkali,  water  is  formed 
and  the  acid  carbonate  (bicarbonate)  remains  dissolved: 

H2CO3  +  NaOH  fc»  H2O  +  NaHCO3,  or  H'  +  OH7  -»  H2O. 

Although  the  bicarbonate  is  technically  an  acid  salt,  its  solution  is 
neutral  on  account  of  the  exceedingly  slight  dissociation  of  the  HCO/ 


324  COLLEGE   CHEMISTRY 

ion.  By  addition  of  an  equivalent  of  sodium  hydroxide  the  normal 
carbonate  is  obtained: 

NaOH  +  NaHC03^>H20  +  Na2CO3,    or    OH'+HC03^H2O  +  C03". 

This  solution,  like  that  of  all  salts  of  a  strong  base  and  a  feeble  acid 
(cf.  p.  254),  is  alkaline  in  reaction.  This  is  because  the  tendency  to 
form  the  very  slightly  ionized  HCO3'  makes  the  foregoing  ionic 
action  noticeably  reversible  (cf.  pp.  254,  311). 

The  normal  carbonates,  with  the  exception  of  those  of  potassium, 
sodium,  and  ammonium,  are  insoluble  in  water,  and  may  be  obtained 
by  precipitation  when  the  proper  ions  are  employed.  For  example  : 

+  2NaCl    or    B 


The  aqueous  solution  of  carbon  dioxide  interacts  with  solutions  of 
barium  and  calcium  hydroxides  in  a  similar  manner: 

Ca(OH)2  +  H,C03  <=?  CaC03  J,  +  H,O. 

These  precipitations  are  used  as  tests  for  carbon  dioxide. 

Excess  of  carbon  dioxide  converts  calcium  carbonate  into  the  more 
soluble  bicarbonate,  and  hence  considerable  quantities  of  "  lime  " 
are  frequently  held  in  solution  by  natural  waters: 

H2CO3  +  CaCO3  ?±  H2Ca(C03)2. 

In  the  same  fashion,  the  carbonates  of  iron  (FeC03),  magnesium, 
and  zinc  are  somewhat  soluble  in  water  containing  free  carbonic  acid. 
In  fact,  the  solution,  transportation,  and  deposition  of  all  these  car- 
bonates take  place  in  nature  on  a  large  scale  by  the  alternate  progress 
and  reversal  of  this  action.  Water  containing  "  lime  "  in  solution 
is  known  as  hard  water  (q.v.). 

Mole  of  Chlorophyll-bearing  Plants  in  Storing  Energy.  — 

While  in  plants  the  same  consumption  of  oxygen  and  production  of 
carbon  dioxide  goes  on  as  in  animals,  only  with  less  rapidity,  an 
action  which  is  in  a  general  way  a  reversal  of  this  takes  place  at 
the  same  time.  The  chlorophyll  and  protoplasm  in  the  leaves  of  the 
plant  have  the  power  of  taking  up  carbon  dioxide.  Part  of  the 
oxygen  is  restored  to  the  air,  and  the  rest  of  the  substance,  including 
all  the  carbon,  is  used  by  the  plant  as  food.  This  operation  goes  on 
only  in  sunlight.  The  various  chemical  compounds  which  plants 
construct  in  large  quantities,  of  which  sugar,  starch,  and  cellulose 


CARBON    AND   THE    OXIDES    OF   CARBON  325 

are  prominent  examples,  are  built  up  as  the  result  of  this  action. 
We  may  represent  the  production  of  cellulose  by  means  of  the 
equation : 

6CO2  +  5H2O  ->  C6H10O5  +  6O,  -  671,000  calories. 

The  energy  absorbed  in  this  action  is  furnished  by  the  sunlight. 
The  same  amount  is  recovered  when  the  wood  is  burned,  and  the 
action  reversed. 

CARBON  MONOXIDE  CO. 

Preparation. —  Carbon  monoxide  is  formed  in  many  industrial 
operations.  We  commonly  observe  the  blue  flame  of  burning  carbon 
monoxide  playing  on  the  surface  of  a  coal  fire.  The  gas  is  produced 
by  the  passage  of  the  carbon  dioxide,  which  is  first  formed,  through 
the  upper  layers  of  heated  coal: 

CO3  +  C  ->  2CO. 

A  similar  reduction  of  carbon  dioxide  is  produced  by  metals  such  as 
zinc,  when  a  moderate  heat  is  applied: 

CO2  +  Zn  ->  ZnO  +  CO. 

On  a  large  scale,  a  mixture  of  carbon  monoxide  and  hydrogen  is 
known  as  water  gas.     Steam  is  turned  into  an  iron  cylinder  lined 
with  fire  clay  and  filled  with  vigorously  burning  coke: 
C  +  H2O  ->  CO  +  H2. 

The  products  are  both  combustible,  and,  by  the  addition  of  sub- 
stances which  burn  with  a  luminous  flame,  the  mixture  is  used  for 
the  manufacture  of  illuminating-gas  (q.v.). 

In  the  laboratory,  carbon  monoxide  is  frequently  obtained  by 
heating  oxalic  acid,  a  solid,  white,  crystalline  substance,  with  con- 
centrated sulphuric  acid.  The  latter  is  here  employed  simply  as  a 
dehydrating  agent  (p.  263),  so  that  it  need  not  be  included  in  the 

equation:  H2C2O4  -  CO2  +  CO  +  H2O. 

To  obtain  pure  carbon  monoxide  from  this  mixture  it  is  necessary  to 
remove  the  carbon  dioxide  by  passing  the  gas  through  a  solution  of 
potassium  hydroxide.     By  using  formic  acid,  or  sodium  formate, 
with  sulphuric  acid,  the  presence  of  the  carbon  dioxide  is  avoided: 
HCHO2  ->  CO  +  H20. 


326  COLLEGE   CHEMISTRY 

Physical  Properties. — Carbon  monoxide  is  a  colorless,  tasteless, 
odorless  gas.  It  is  very  slightly  soluble  in  water.  Its  density  is 
almost  the  same  as  that  of  air.  When  liquefied  it  boils  at  —  190°. 

Chemical  Properties.  —  All  the  chemical  properties  of  carbon 
monoxide  are  referable  to  the  fact  that  in  it  the  element  carbon 
appears  to  be  bivalent:  C  =  O.  The  compound  is  in  fact  unsatu- 
rated,  and  combines  with  oxygen,  chlorine,  and  other  substances 
directly.  Thus  the  gas  burns  in  the  air,  uniting  with  oxygen  to  form 
carbon  dioxide.  Again,  iron  (q.v.)  is  manufactured  by  the  reduction 
of  the  oxide  of  iron  by  gaseous  carbon  monoxide  in  the  blast  furnace: 
Fe2O3  +  SCO  <=>  2Fe  +  3CO2. 

In  sunlight  carbon  monoxide  unites  directly  with  chlorine  to  form 
carbonyl  chloride  COC12.  It  unites  directly  with  nickel  and  iron  to 
form  nickel  carbonyl  and  iron  carbonyl  (q.v.),  respectively. 

The  gas  is  an  active  poison.  When  inhaled  it  unites  with  the 
haemoglobin  of  the  blood  to  the  exclusion  of  the  oxygen  which  forms 
a  less  stable  compound  (cf.  p.  52).  A  quantity  equivalent  to  about 
10  c.c.  of  the  gas  per  kilo,  weight  of  the  animal  is  sufficient  to  pro- 
duce death,  about  one-third  of  the  whole  haemoglobin  having  entered 
permanently  into  combination  with  carbon  monoxide. 

Carbon  Disulphide  CS2.  —  This  compound  is  made  by  direct 
union  of  sulphur  vapor  and  glowing  charcoal.  An  electric  furnace 
like  that  in  Fig.  57  (p.  304)  is  employed.  The  substance  comes  off 
as  a  vapor  and  is  condensed. 

Carbon  disulphide  is  a  colorless,  highly  refracting  liquid  (b.-p.  46°). 
Traces  of  .other  compounds  give  the  commercial  article  a  disagreeable 
smell.  It  burns  in  air,  forming  carbon  dioxide  and  sulphur  dioxide. 
Iodine,  phosphorus,  sulphur,  and  rubber  dissolve  freely  in  it. 

Exercises.  —  1.  To  which  of  the  factors  in  the  interaction  of  cal- 
cium carbonate  and  hydrochloric  acid  (p.  322)  is  due  the  forward 
displacement  of  all  the  equilibria? 

2.  What  will  be  the  excess  of  pressure  inside  a  bottle  of  soda- 
water  when  four  volumes  of  carbon  dioxide  are  dissolved  in  one 
volume  of  water? 

3.  What  volume  of  liquid  carbon  dioxide,  measured  at  0°,  will  be 
required  to  give  75  liters  of  the  gas  at  0°  and  760  mm.  pressure? 

4.  What  are  the  exact  relative  weights  of  equal  volumes  of  carbon 
dioxide,  carbon  monoxide,  air,  and  steam? 


I 

CHAPTER  XXIX 
SOME    CARBON    COMPOUNDS 

...  THE  compounds  of  carbon  with  hydrogen  are  called  hydrocarbons. 
Those  containing  oxygen,  as  well,  are  divided  into  numerous  and 
extensive  groups  according  to  their  behavior.  Thus  there  are  acids 
like  acetic  acid,  carbohydrates  like  sugar  and  starch,  alcohols  like 
common  (ethyl)  alcohol,  and  esters  like  ethyl  acetate  and  fat.  There 
are  also  compounds  related  to  cyanogen,  like  prussic  acid,  which  con- 
tain nitrogen.  We  can  discuss  only  one  or  two  examples  from  each 
of  the  groups  named. 

THE  HYDROCARBONS. 

More  than  two  hundred  and  fifty  compounds  of  carbon  and  hydro- 
gen have  been  described.  They  fall  into  several  distinct  series,  the 
chief  one  of  which  contains  methane  CH4  as  its  simplest  member. 
On  account  of  the  fact  that  certain  members  of  this  set  are  found  in 
paraffin,  it  is  commonly  known  as  the  paraffin  series.  For  the  reason 
that  in  this  series  the  carbon  has  all  its  four  valences  employed,  the* 
members  are  also  called  the  saturated  hydrocarbons. 

Paraffin  Series  of  Hydrocarbon  ft.  —  The  names  and  formulae 

of  some  of  the  members  of  the  series  are  : 

Gases:  methane  CH4,  ethane  C2H6,  propane  C3H8,  butane  C4H,0  (b.-p.  1°). 

Liquids:  pentane  C5H12  (b.-p.  35°),  hexane  C^H.l4  (b.-p.  71°),  heptane  C7H]6 
(b.-p.  99°),  ---- 

Solids:  hexadecane  C^H^  (m.-p.  18°),  pentatriacontane  C^H^  (m.-p.  74.7°), 
hexacontane  C^H^s  (m.-p.  102°). 


In  composition,  each  is  related  to  the  preceding  one  by  containing  the 
additional  units  CH2.  Their  relations  will  be  more  clearly  perceived 
if  we  employ  the  graphic  formulae  for  the  first  three  members.  The 
hydrogen  is  univalent,  and  the  carbon  quadrivalent: 

327 


328 


COLLEGE   CHEMISTRY 


H 

I 
H  -  C-  H 

I 
H 


H      H 

I        I 

H  -  C  -  C  -  H 

I        I 
H     H 


H     H     H 

I        I        I 
H-C-C-C-H 

I        I       I 
H     H     H 


The  natural  gas  of  Pennsylvania  and  Ohio  contains  a  large  propor- 
tion of  methane. 

Petroleum  consists  of  a  mixture  of  the  liquid  and  solid  members  of 
the  series  in  varying  proportions,  and  is  found  in  many  parts  of  the 
United  States,  in  Ontario,  at  Baku  on  the  Caspian,  in  India,  and  in 
Japan.  In  oil-refining,  advantage  is  taken  of  the  differences  in  the 
boiling-points  to  make  a  partial  separation  of  the  components  by 
fractional  distillation.  In  this  operation  a  thermometer  immersed  in 
the  vapor  (Fig.  14,  p.  26)  rises  steadily  as  the  less  volatile  compo- 
nents increase  in  quantity  and  the  more  volatile  decrease.  The 
receiver  is  changed  ^as  certain  temperatures  are  reached. 


Name. 

Components. 

B.-P. 

Uses. 

Petroleum  ether 
Gasolene  .    .    . 
Naphtha  .    .    . 
Benzine    .    .    . 
Kerosene     .    . 

Pentane,  hexane 
Hexaiie,  heptane 
Heptane,  octane 
Octane,  nonane 
Decane,  hexadecane 

40°  -  70° 
70°  -  90° 
80°  -120° 
120°  -150° 
150°  -300° 

Solvent,  gas-making 

fuel 
« 

Illuminating  -oil 

The  portions  of  still  higher  boiling-point  are  used  as  lubricating 
oils.  Vaseline  consists  of  substances  C22H46  to  CjgH^  (m.-p.  40-50°). 

By  cooling  the  residues  from  the  retorts  with  a  freezing  mixture 
(cf.  p.  204),  some  of  the  solid  members  of  the  series,  C^H^  to  C28H58, 
are  obtained  as  white  flakes,  which  are  separated  by  filtration  in 
presses.  This  material  forms  the  paraffin  used  in  waterproofing 
paper,  in  laundry  work,  and  as  an  ingredient  in  candles. 

The  hydrocarbons  are  extremely  indifferent  in  their  chemical  be- 
havior. They  have  none  of  the  properties  of  acids,  bases,  or  salts. 
The  halogens,  notably  chlorine  and  bromine,  however,  interact  with 
them  (see  below).  When  burned  they  all  produce  carbon  dioxide 
and  water. 

Methane  CH4.  —  Methane,  otherwise  known  as  marsh-gas,  is 
the  chief  component  of  natural  gas.  It  also  rises  to  the  surface  when 


THE    HYDROCARBONS  329 

the  bottoms  of  marshy  pools  are  disturbed,  and  issues  from  seams  in 
coal  beds  ("  fire-damp  "  from  Ger.  Dampf,  vapor).  In  these  two 
cases  it  results  from  the  decomposition  of  vegetable  matter  in  absence 
of  air. 

Methane  may  be  made  from  inorganic  materials  by  the  action  of 
water  upon  aluminium  carbide,  prepared  by  the  interaction  of  alumi- 
nium oxide  and  carbon  in  the  electric  furnace  (cf.  pp.  304,  321) : 
A14C3  +  12H2O-*4A1(OH)3  +  3CH4. 

In  the  laboratory  the  gas  is  commonly  obtained  by  the  distillation 
of  a  dry  mixture  of  sodium  acetate  and  sodium  hydroxide: 
NaC2H3O2  +  NaOH  -»  Na2CO3  +  CH4. 

When  a  mixture  of  methane  and  chlorine  is  exposed  to  sunlight 
several  changes  occur  in  succession  (cf.  p.  115): 

CH4  +  C12  -*  CH3C1  +  HC1,  CH3C1  +  C12  -» CH2C12  +  HC1, 
CH2C12  +  C12  ->  CHC13  +  HC1,  CHC13  +  C12  -*  CC14  +  HC1. 
This  kind  of  interaction  with  the  halogens  is  characteristic  of  com- 
pounds of  hydrogen  and  carbon.  It  takes  place  slowly,  and  is  there- 
fore entirely  different  from  ionic  chemical  change.  It  consists  in  a 
progressive  substitution  of  chlorine  for  hydrogen,  unit  by  unit. 
Chloroform  CHC13  is  the  only  one  of  the  four  products  which  is  a 
familiar  substance.  The  corresponding  iodine  derivative,  iodoform 
CHI3  is  used  in  surgical  dressing.  These  substances  are  not  salts, 
and  are  not  ionized  in  solution.  They  are  very  slowly  hydrolyzed 
by  water,  —  carbon  tetrachloride,  for  example,  giving  carbonic  acid 
and  hydrochloric  acid. 

Ethylene  C2H4.  —  Ethylene  is  the  first  member  of  the  second 

series  of  hydrocarbons.     It  corresponds  to  ethane  C2H6,  but  contains 

in  each  molecule  two  hydrogen  units  less  than  does  this  substance. 

Ethylene  is   made  by  heating  common  alcohol  (ethyl  alcohol) 

with  concentrated  sulphuric  acid: 

C2H5OH-»H2O  +  C3H4f. 

A  comparison  of  the  graphic  formulae  of  the  alcohol  and  ethylene 
shows  that  this  loss  of  water  leaves  the  carbon  partly  unsaturated: 
H    H  H    H  H      H 

l_| JJ_ 

H    H 


330  COLLEGE   CHEMISTRY 

The  water  may  also  be  removed  by  allowing  alcohol  to  fall  drop  by 
drop  into  heated  phosphoric  acid. 

Ethylene  is  formed,  along  with  acetylene  and  other  substances, 
when  any  saturated  hydrocarbon  is  heated  strongly.     Even  methane 

givesit: 


Ethylene  is  a  gas.  It  burns  in  the  air  with  a  flame  which,  on 
account  of  the  great  separation  of  free  carbon  which  takes  place 
temporarily  during  the  combustion  (cf.  Flame),  is  highly  luminous. 
It  will  be  seen  that,  in  the  formula,  but  three  of  the  valences  of  each 
carbon  unit  are  occupied:  the  substance  is  unsaturated.  Hence, 
when  ethylene  is  passed  through  liquid  bromine  it  is  rapidly  absorbed, 
and  the  bromine  seems  to  increase  in  volume  and  finally  loses  all  its 
color,  being  converted  into  a  transparent  liquid  having  the  com- 
position C2H4Br2,  ethylene  bromide. 

Acetylene.  —  This  substance,  likewise  a  gas,  is  the  first  member  of 
still  another  unsaturated  series.  Its  formula  C2H2  shows  that  its 
molecule  lacks  four  of  the  hydrogen  units  necessary  to  the  complete 
saturation  which  we  find  in  ethane.  Graphically  its  structure  is 
usually  represented  thus  :  H  —  C  =  C  —  H.  This  gas  is  formed  in 
small  quantities  by  direct  union  of  carbon  and  hydrogen  in  the 
electric  arc  (p.  320).  It  is  also  produced  when  ethylene  is  passed 
through  a  heated  tube:  C2H4  —  >  C2H2  -f  H2  (cf.  Flame).  When  cal- 
cium carbide  (p.  321)  is  thrown  into  water  it  is  hydrolyzed.  Violent 
effervescence  occurs,  the  calcium  carbide  is  disintegrated,  a  precipi- 
tate of  calcium  hydroxide  is  formed,  and  acetylene  passes  off  as  a 

CaC2  +  2H2O  ->  Ca(OH)2  +  C,!^ 

This  action  is  like  that  of  water  on  calcium  phosphide  (p.  307), 
and  magnesium  nitride  (p.  281). 

Acetylene  burns  with  a  flame  which  is  still  more  luminous  than 
that  of  ethylene. 

THE  ACIDS. 

Formic  Acid  HCO2H.  —  The  removal  of  water  from  formic 
acid  produces  carbon  monoxide  (p.  325).  Although  we  cannot 
reverse  the  process  and  cause  carbon  monoxide  to  combine  with 
water,  we  can  make  it  unite  with  bases.  By  passing  carbon  mon- 
oxide over  hot  sodium  hydroxide,  we  obtain  sodium  formate,  from 
I 


THE   ACIDS  331 

which  formic  acid  may  be  liberated  by  double  decomposition  with 
another  acid:  CQ  +  NaQH  _^  NaCOjH 

This  acid  is  secreted  by  red  ants,  and  is  found  in  stinging  nettles. 
It  is  a  liquid  boiling  at  100. 1°.  The  molecule  contains  two  atoms  of 
hydrogen,  but  the  acid  is,  in  fact,  monobasic. 

Acetic  Acid  JEfC2H3O2  —  This  acid  is  produced  in  the  dry  dis- 
tillation of  wood*  (p.  318).  Large  quantities  of  it  are  manufactured 
from  dilute  alcohol.  The  liquid  is  allowed  to  flow  in  a  slow  stream 
through  a  barrel  filled  with  shavings.  Holes  in  the  barrel  provide 
for  the  access  of  air,  and  a  bacterium  with  which  the  shavings  are 
infected  promotes  oxidation  of  the  alcohol: 

C2H5OH  +  O2  ->  C2H3OOH  +  H20. 

Oxygen  alone  does  not  affect  alcohol  in  the  cold.  The  bacterium  (B. 
aceti,  "  mother-of- vinegar  ")  assists  this  action,  as  lower  organisms 
are  found  to  assist  many  chemical  actions,  in  a  way  which  may  be 
described  roughly  as  catalytic  (see  Fermentation,  p.  332). 

The  dilute  solution  of  acetic  acid  produced  in  this  manner  contains 
from  five  to  thirteen  per  cent  of  acetic  acid,  and  is  known  as  vinegar. 
By  fractional  distillation  the  solution  may  be  concentrated  until  a 
little  water  only  remains,  and  finally,  by  freezing  (cf.  p.  204),  the 
acetic  acid  may  be  crystallized  out.  Pure  acetic  acid,  in  consequence 
of  its  freezing  readily  in  cold  weather,  is  known  as  "  glacial  "  acetic 
acid.  It  melts  at  16.7°  and  boils  at  118°. 

Although  four  atoms  of  hydrogen  are  contained  in  its  molecule, 
but  one  of  these  is  replaceable  by  metals.  This  fact  is  recognized  in 
the  reaction  formula  (p.  83)  of  the  acid,  HC2H302. 

Oxalic  Acid  H2C2O4.  —  This  acid  is  dibasic.  Its  calcium  salt 
CaC204  is  the  least  soluble  of  the  salts  of  calcium,  and  is  found  in 
many  plants  in  the  form  of  bundles  of  needle-shaped  crystals.  The 
acid  may  be  made  by  oxidation  of  sugar  with  nitric  acid.  The  white 
crystalline  hydrate,  H2C204,2H2O,  is  the  form  used  in  the  laboratory. 

*  The  dry  distillation  of  bones  (p.  318),  on  the  other  hand,  and  of  animal 
matter  (p.  281)  in  general,  gives  alkaline  liquids,  because  of  the  ammonia 
that  is  formed. 


332  COLLEGE  CHEMISTRY 

CARBOHYDRATES  AND  FERMENTATION. 

Carbohydrates.  —  The  various  kinds  of  sugar,  starch,  and  cellu- 
lose form  a  closely  related  group  of  substances.  As  it  happens  that 
the  proportion  of  hydrogen  to  oxygen  in  the  composition  of  most  of 
them  is  the  same  as  that  of  these  elements  in  water,  they  are  known 
by  the  name  of  carbohydrates.  None  of  these  substances  show  any 
distinct  evidence  of  ionization. 

Dextrose,  otherwise  known  as  glucose  or  grape-sugar,  is  a  white 
crystalline  substance  having  the  composition  C6H12O6.  It  is  found 
dissolved  in  the  juices  of  sweet  fruits,  such  as  grapes. 

The  most  familiar  sugar  is  cane-sugar  C12H22OU,  which  may  form 
as  much  as  18  per  cent  by  weight  of  the  juices  pressed  from  the 
sugar-cane,  and  sometimes  reaches  16  per  cent  of  the  fluid  material 
in  the  sugar-beet.  It  is  prepared  by  boiling  the  juices  with  animal 
charcoal  (p.  319),  to  remove  the  coloring  matter  which  would  other- 
wise give  the  sugar  a  brown  tint.  The  liquid  is  then  concentrated 
until  crystals  appear.  The  mother-liquor  which  no  longer  deposits 
crystals  is  known  as  molasses. 

The  relation  of  starch  to  the  sugars  is  seen,  not  only  in  the  formula 
(C6H1005)n,  but  in  the  fact  that  by  boiling  starch  with  dilute  acids, 
dextrose  is  formed  along  with  other  products  of  the  hydrolysis. 
Commercial  "  glucose  "  is  made  by  this  process.  Starch  is  an  in- 
soluble white  substance  which  is  found  in  the  form  of  fine  particles 
in  the  fruit  and  other  parts  of  plants. 

Cellulose  has  the  same  composition  as  starch.  It  forms  the  frame- 
work of  the  cells  of  plants.  In  many  cases  it  is  overlaid  with  a  con- 
siderable thickness  of  lignin,  which  in  paper-making  is  removed  by 
boiling  the  wood  with  sodium  hydroxide  or  calcium  bisulphite  solu- 
tion (p.  266).  When  the  product  has  been  washed  thoroughly  with 
water,  almost  pure  cellulose  remains.  Matted  cellulose  in  thin 
sheets  forms  the  basis  of  paper,  and  filter  paper  contains  nothing  else 
(see  under  Aluminium  sulphate).  Other  forms  of  pure  cellulose  are 
known  as  cotton,  linen,  and  jute,  according  to  their  sources.  Al- 
though the  chemical  composition  of  these  varieties  is  identical,  the 
physical  properties  vary  considerably. 

Fermentation.  —  This  is  the  name  given  to  a  number  of  different 
chemical  changes,  brought  about  by  catalytic  action  of  complex 


ALCOHOLS,   ESTERS,    AND   SOAP 

chemical  compounds  secreted  by  living  organisms.  These  com- 
pounds are  called  enzymes,  and,  in  many  cases,  have  been  separated 
from  the  organisms  by  means  of  solvents.  Their  action  must  be 
regarded  as  catalytic,  since  small  quantities  of  the  active  organisms 
or  of  the  enzymes  can  produce  very  extensive  chemical  changes 
without  themselves  suffering  alteration  in  the  process. 

The  organisms  may  be  divided  into  three  classes,  each  secreting 
different  enzymes  which  confine  themselves  for  the  most  part  to 
special  kinds  of  chemical  change.  (1)  The  molds,  when  grown  in 
sugar  solution  or  beef  extract,  or  other  nutritive  solutions,  produce 
decompositions  known  collectively  as  putrefaction.  (2)  Certain 
bacteria  promote  the  oxidation  of  alcohol  to  acetic  acid  (p.  331). 
Some  also  decompose  sugar,  furnishing  butyric  or  lactic  acid  as  one 
of  the  products.  (3)  The  yeasts  (saccharomycetes)  flourish  in  solu- 
tions of  some  sugars,  and  decompose  them  into  alcohol  and  carbon 
dioxide.  This  decomposition  is  known  as  alcoholic  fermentation. 
These  changes  are  usually  brought  about  by  actual  introduction  of 
the  organism.  In  brewing,  however,  the  enzyme  itself,  diastase,  is 
employed  to  hydrolyze  starch. 

When  yeast  is  added  to  a  solution  of  cane-sugar,  hydrolysis  into 
dextrose  and  levulose  first  occurs: 

C12H22On  +  H20  ->  C6H1200  +  C6H  1206. 

This  change  is  produced  also  by  boiling  with  dilute  acids.  But 
the  enzyme  of  the  yeast  immediately  decomposes  the  dextrose  and 
the  levulose  into  alcohol  and  carbon  dioxide: 

C6H1206  -»  2C2H5OH  -f  2C02. 
The  liquid  effervesces,  and  the  carbon  dioxide  escapes  into  the  air. 

ALCOHOLS,  ESTERS,  AND  SOAP. 

Alcohols.—  When  wood  is  distilled  (p.  3 18),  methyl  alcohol  CH3OH 
is  found  in  the  fluid  product.  When  purified  this  is  a  colorless 
liquid  boiling  at  66°.  Its  solution  in  water  shows  no  evidence  of 
ionization. 

Common  or  ethyl  alcohol  C2H5OH  is  formed  in  the  fermentation  of 
solutions  of  sugar  by  yeast  (p.  333),  and  is  separated  from  the  water 
and  the  other  products  of  fermentation  by  distillation.  The  product 
contains  95%  of  alcohol  by  volume  (in  Great  Britain,  90%)  and  is 


334  COLLEGE    CHEMISTRY 

applicable  to  most  commercial  uses.  Absolute  alcohol,  entirely  free 
from  water,  is  made  by  placing  the  spirit  in  vessels  filled  with 
quicklime.  The  latter  interacts  with  the  water  producing  calcium 
hydroxide,  and  the  clear  liquid  which  is  poured  off  is  distilled  once 
more.  Pure  alcohol  boils  at  78.3°. 

Glycerine  C3H5(OH)3  is  an  alcohol  containing  three  hydroxyl 
groups,  —  a  trihydric  alcohol. 

Esters:  Soap.  —  When  an  acid  and  an  alcohol  are  mixed,  an 
ester  and  water  are  formed.  The  action  is  slow  and,  being  reversible, 
is  always  incomplete.  But  by  introduction  of  a  dehydrating  agent, 
like  concentrated  sulphuric  acid,  the  water  is  removed  and  the 
change  brought  to  completion.  Thus,  ethyl  alcohol  and  acetic 
acid,  when  warmed  with  sulphuric  acid,  give  ethyl  acetate: 

C2H5OH  +  HC2H3O2  <=>  C2H5C2H3O2  +  H2O. 

When  this,  or  any  other  ester,  is  boiled  with  excess  of  water,  the 
foregoing  action  is  reversed. 

When  esters  are  boiled  with  strong  bases,  such  as  sodium  hydroxide 
solution,  an  alcohol  and  the  salt  of  the  acid  are  formed : 

NaOH  +  C2H5C2H3O2  ->  C2H5OH  +  NaC2H3O2. 

The  fats  are  mixtures  of  complex  esters  and  undergo  this  change 
just  as  do  the  simpler  esters.  The  sodium  salts  produced  from  fats 
are  known  as  soaps,  and  this  general  kind  of  action  is  called,  therefore, 
saponification  (Lat.  sapo,  soap). 

In  beef  fat  the  chief  esters  present  are  tripalmitin,  tristearin,  and 
triolein.  It  will  be  sufficient  to  illustrate  the  chemistry  of  the  manu- 
facture of  soap  by  discussing  the  case  of  one  of  these  substances. 
Tripalmitin  is  the  glyceryl  ester  of  palmitic  acid.  When  fat  is  mixed 
with  hot  sodium  hydroxide  solution  it  first  forms  an  emulsion*  in 
which  the  fat  is  disseminated  in  minute  droplets  through  the  liquid. 
This  is  a  result  of  surface  tension.  When  the  emulsion  is  boiled, 
the  fat  is  slowly  decomposed  into  glycerine  and  sodium  palmitate. 
The  change  is  similar  in  plan  to  the  simpler  one  just  discussed: 

3NaOH  +  C3H5(C16H3102)3->C3H&(OH)3  +  3NaC1BH31O2. 

*  In  the  intestines  the  same  office  is  performed  by  the  gall,  secreted  by  the 
liver,  and  so  the  fat  is  prepared  for  absorption  into  the  system. 


CYANOGEN  335 

Changes  similar  to  this  occur  with  the  other  two  esters.  The  only 
difference  is  that  the  organic  radical  in  the  case  of  tristearin  is 
C^H^O-j,  and  in  the  case  of  triolein  C18H3302.  Both  products  in  each 
of  the  three  cases  are  soluble  in  water,  but  when  common  salt  is 
added  to  the  solution  the  sodium  salts  of  the  organic  acids  are  separ- 
ated ("  salted  out  ")  as  a  solid  mass,  which  is  known  as  soap.  When 
potassium  hydroxide  takes  the  place  of  sodium  hydroxide  the  whole 
unsalted  mass  is  semi-fluid,  and  is  known  as  soft  soap. 

The  components  of  soap,  like  other  soluble  salts,  are  highly  ionized 
in  solution,  and  show  all  the  properties  of  ionogens.  For  example, 
addition  of  an  acid  liberates  the  organic  acid.  So  also,  when  soap  is 
dissolved  in  hard  water  (cf.  p.  77),  a  white,  flocculent  precipitate  is 
formed,  which  coagulates  upon  the  sides  of  the  vessel.  This  is  a 
mixture  of  the  calcium  salts  formed  by  union  of  the  proper  ions. 
For  example,  the  sodium  palmitate  is  changed  as  follows: 

2NaC16H31O2  +  CaS04->Ca(C15H31COO)2|+  Na2SO4. 

Most  of  the  salts  of  these  acids,  with  the  exception  of  those  of  potas- 
sium and  sodium,  are  insoluble  in  water. 

CYANOGEN. 

Cyanogen  C2N2.  —  This  compound  is  prepared  by  allowing  a 
solution  of  cupric  sulphate  to  trickle  into  a  warm  solution  of  potas- 
sium cyanide.  The  cupric  cyanide,  at  first  precipitated,  quickly 
decomposes,  giving  cuprous  cyanide  and  cyanogen: 

2KNC  +  CuSO4-*Cu(NC)2|  +  K2SO4, 

2Cu(NC)2  -»  2CuNC  +  C2N2f 

Cyanogen  is  a  very  poisonous  gas  with  a  characteristic,  faint  odor. 

Hydrocyanic  Acid  HNC.  —  This  acid,  called  also  prussic  acid, 

is  most  easily  made  by  the  action  of  an  acid  upon  a  cyanide  (see 
Potassium  cyanide)  followed  by  distillation.  It  is  a  colorless  liquid 
boiling  at  26.5°.  It  has  an  odor  like  that  of  bitter  almonds,  and  is 
highly  poisonous.  In  aqueous  solution  it  is  an  extremely  feeble 
acid.  Hydrocyanic  acid  and  the  cyanides  are  unsaturated  com- 
pounds, a  fact  which  is  illustrated  in  the  two  following  paragraphs. 

Cyanates  and  Thiocyanates.  —  When  potassium  cyanide  is 
fused  and  stirred  with  an  easily  reducible  oxide,  like  lead  oxide 


336  COLLEGE    CHEMISTRY 

(PbO),  the  metal  (for  example,  the  lead)  collects  at  the  bottom  of 
the  iron  crucible  in  molten  form,  and  potassium  cyanate  KNCO  is 
produced:  KNQ  +  pbQ  _^  KNCQ  +  pb 

When  potassium  cyanide  in  aqueous  solution  is  boiled  with  sulphur 
or  with  a  polysulphide  (p.  255),  it  is  converted  into  potassium  thio- 
cyanate  KCNS.  This  salt,  or  ammonium  thiocyanate  NH4CNS,  is  used 
in  testing  for  ferric  ions  on  account  of  the  deep-red  color  of  ferric 
thiocyanate  (cf.  p.  179). 

Exercises. —  1.  Make  the  graphic  formulae  of  hexane  (p.  327), 
ethyl  formate,  ethylene  bromide  (p.  330),  ethyl  alcohol. 

2.  Make  equations  for  the  formation  of  aluminium  carbide  (p.  329) , 
the  saponification  of  triolein  (p.  334). 

3.  Prepare  a  summary  of  the  various  statements  that  have  been 
made  in  the  text  about  catalysis  (e.g.  pp.  54,  74,  110,  114,  161,  257, 
268,  302,  333),  and  illustrate  fully. 


CHAPTER  XXX 


FLAME    AND    ILLUMINANTS 

Meaning  of  the  Term.  —  In  the  combustion  of  charcoal  there 
is  hardly  any  flame,  for  the  light  emanates  almost  entirely  from  the 
incandescent,  massive  solid.     When  two  gases  are  mixed  and  set  on 
fire,  a  sort  of  flame  passes  through  the 
mixture,  but  this  can  hardly  be  accounted 
a   flame,  in   the   ordinary   sense,    either. 
The  rapid  movement  of  the  flash,  and  the 
explosion  which  accompanies  it,  are  in  a 
manner  the  precise  opposite  of  the  quiet 
combustion    which    is    characteristic    of 
flames. 

With  illuminating-gas  the  production 
of  its  very  characteristic  flame  is  due  to 
the  chemical  union  of  a  stream  of  one 
kind  of  gas  in  an  atmosphere  of  another. 
The  flame  is  made  up  of  the  heated  matter 
where  the  two  gases  meet.  In  the  case  of 
a  burning  candle,  one  of  the  bodies 
appears  to  be  a  solid,  but  a  closer  scrutiny 
of  the  phenomenon  shows  that  the  solid 
does  not  burn  directly.  A  combustible 
gas  is  manufactured  continuously  by  the 
heat  of  the  combustion  and  rises  from  the 
wick.  The  introduction  of  a  narrow  tube 
into  the  interior  of  the  flame  enables  us 
to  draw  off  a  stream  of  this  gas  and  to 
ignite  it  at  a  remote  point.  Thus,  a  flame 

is  a  phenomenon  produced  at  the  surface  where  two  gases  meet  and 
undergo  combination  with  the  evolution  of  heat  and,  more  or  less,  light. 

In  the  chemical  point  of  view,  it  is  a  matter  of  indifference  whether 
the  gas  outside  the  flame  contains  oxygen,  and  the  gas  inside  consists 

337 


FIG.  60. 


338  COLLEGE   CHEMISTRY 

of  substances  ordinarily  known  as  combustibles,  or  whether  this 
order  is  reversed.  In  an  atmosphere  of  ordinary  illuminating-gas, 
the  flame  must  be  fed  with  air.  This  condition  is  easily  realized 
(Fig.  60) .  The  lamp-chimney  is  closed  at  the  top  until  it  has  become 
filled  with  illuminating-gas.  After  the  lapse  of  a  few  minutes  this 
can  be  ignited  as  it  issues  from  the  bottom  of  the  wide,  straight  tube 
which  projects  from  the  interior.  When  the  hole  in  the  cover  of 
the  lamp-chimney  is  then  opened,  the  upward  draft  causes  the 
flame  of  the  burning  gas  to  recede  up  the  tube,  and  there  results 
a  flame  fed  by  air  and  burning  in  coal-gas.  In  an  atmosphere 
of  this  kind,  materials  playing  the  part  of  a  candle  burning  in 
air  would  have  to  be  substances  which,  under  the  influence  of  the 
heat  of  combustion,  give  off  oxygen.  Strongly  heated  potassium 
chlorate,  for  example,  appears  to  burn  continuously  in  such  an 
atmosphere  when  lowered  into  it  in  a  deflagrating  spoon. 

Luminous  Flames.  —  The  flame  of  hydrogen,  under  ordinary 
circumstances,  is  almost  invisible,  nearly  all  the  energy  of  the  com- 
bustion being  devoted  to  the  production  of  heat.  A  part  of  this, 
however,  may  be  transformed  into  light  by  the  suspension  of  a 
suitable  solid  body,  such  as  a  platinum  wire,  in  the  flame.  The 
holding  of  a  piece  of  quicklime  in  an  oxyhydrogen  flame  (cf.  p.  74) 
is  a  practical  illustration  of  this  method  of  securing  luminosity. 
In  general,  luminosity  may  be  produced  by  the  presence  of  some 
incandescent  solid. 

In  the  Welsbach  lamp  the  flame  itself  is  non-luminous,  and,  but 
for  the  mantle,  would  be  identical  with  the  ordinary  Bunsen  flame. 
The  mantle  which  hangs  in  the  flame,  however,  by  its  incandescence, 
furnishes  the  light.  This  mantle  is  composed  of  a  mixture  of  99 
per  cent  thorium  dioxide  (Th02)  and  one  per  cent  cerium  dioxide 
(Ce02).  While  many  cheaper  oxides  would  give  out  a  white  light, 
they  have  neither  sufficient  coherence,  nor  sufficiently  lasting,  high 
emissivity  to  make  their  use  practicable.  It  is  worth  noting  that 
any  appreciable  variation  from  the  above  proportions,  by  the  intro- 
duction of  either  more  or  less  cerium  oxide,  produces  a  marked 
diminution  in  the  intensity  and  whiteness  of  the  light. 

In  cases  of  brilliant  combustion,  as  of  magnesium  ribbon  or  phos- 
phorus, a  solid  body  is  formed  whose  incandescence  accounts  for  the 
light.  The  flame  of  ordinary  illuminating-gas  does  not  at  first  sight 


FLAME   AND   ILLUMINANTS  339 

appear  to  give  evidence  of  the  presence  of  any  solid  body.  But  if  a 
cold  evaporating  dish  is  held  in  the  flame  for  a  moment,  a  thick 
deposit  of  finely  divided  carbon  (soot)  is  formed,  and  we  at  once 
realize  that  the  light  is  due  to  the  glow  of  these  particles  in  a  mass  of 
intensely  hot  gas.  Carbon  is,  indeed,  an  extremely  combustible 
substance,  and  is  eventually  entirely  consumed.  But  a  fresh  supply 
is  continually  being  generated  in  the  interior  of  the  flame,  while  the 
oxygen  with  which  it  is  to  unite  is  outside  the  flame  altogether. 
Thus  the  carbon  particles  persist  until,  drifting  with  the  spreading 
gas,  they  reach  the  periphery  of  the  flame. 

TJie  Bunsen  Flame. —  In  the  burner  devised  by  Robert  Bunsen, 
a  jet  of  ordinary  illuminating-gas  is  projected  from  a  narrow  opening 

into  a  wider  tube.     In  this  tube  it  becomes  mixed  with  air.  drawn 

t 

in  through  openings  whose  dimensions  can  be  altered  by  means  of  a 
perforated  ring.  When  the  supply  of  air  is  sufficient,  the  flame 
becomes  non-luminous.  With  a  somewhat  different  construction, 
and  the  use  of  a  bellows  to  force  a  larger  proportion  of  air  into  the 
gas,  a  still  hotter  flame  can  be  produced.  The  instrument  in  this 
case  is  known  as  a  blast-lamp.  The  high  temperature  of  the  Bunsen 
flame  is  not  difficult  to  account  for.  It  will  be  seen  at  once,  on 
handling  the  burner,  that  the  flame  diminishes  to  one-half  or  one- 
third  of  its  previous  size  when  air  is  admitted.  Since  the  same 
amount  of  gas  is  burning  in  both  cases,  and  the  products  in  both 
cases  are  water  and  carbon  dioxide,  the  total  amounts  of  heat  pro- 
duced must  in  both  cases  be  the  same.  The  production  of  an  equal 
amount  of  heat  in  a  smaller  flame  necessarily  causes  this  to  have  a 
higher  average  temperature. 

Structure  of  the  Illuminating  and  the  Bunsen  Flame.  — 

When  an  exceedingly  small  luminous  flame  is  examined,  the  various 
parts  of  which  it  consists  may  easily  be  made  out.  In  the  interior 
there  is  a  dark  cone  which  is  composed  of  illuminating-gas  and  air, 
and  in  it  no  combustion  is  taking  place.  A  match-head  may  be  held 
here  for  some  time  without  being  set  on  fire.  This  is  therefore  not 
properly  a  part  of  the  flame.  Outside  this  is  a  vivid  blue  layer  (C, 
Fig.  61)  which  is  best  seen  in  the  lower  part  of  the  flame,  but  extends 
beneath  the  luminous  sheath,  and  covers  the  dark  inner  cone  com- 
pletely. Outside  the  blue  flame,  and  covering  the  greater  part  of 


340 


COLLEGE   CHEMISTRY 


it,  is  the  cone-shaped  luminous  portion  (B).  Over  all  is  a  faint 
mantle  of  non-luminous  flame  (A),  which  becomes  visible  when  the 
light  from  the  luminous  part  is  purposely  obstructed.  In  the  lumi- 
nous gas-flame,  therefore,  there  are  four  regions,  if  we  count  the 
inner  cone  of  gas.  The  difference  between  this  and  the 
non-luminous  Bunsen  flame  is  that  in  the  latter  the  lumi- 
nous region  is  omitted,  and  the  inner,  dark  cone,  the  blue 
sheath,  and  the  outer  mantle,  are  the  only  parts  which 
can  be  distinguished. 

Illuminating 'Gas  and  its  Composition.  —  Before 
considering  the  chemistry  of  the  gas-flame,  it  is  neces- 
sary to  know  what  substances  are  burning.  The 
illuminating-gas  in  Europe,  and  in  many  of  the  smaller 
cities  of  the  United  States,  is  usually  coal-gas,  while  in 
the  larger  cities  of  America  it  is  almost  always  made  from 
water-gas.  Coal-gas  is  obtained  by  the  destructive  dis- 
tillation of  soft  coal,  and  is  freed  from  ammonia  (cf.  p. 
281)  and  tar  by  washing  and  cooling,  and  from  hydrogen 
sulphide  and  carbon  dioxide  by  passage  through  layers 
of  slaked  lime.  The  water-gas,  made  by  the  action  of 
steam  upon  anthracite  or  coke,  being  composed  of  car- 
bon monoxide  and  hydrogen  (cf.  p.  325),  has  no  illumi- 
nating power.  It  is  therefore  "carburetted,"  that  is,  mixed  with 
hydrocarbons,  by  passage  through  a  cylindrical  structure  filled  with 
white-hot  firebrick,  upon  which  falls  a  small  stream  of  high-boiling 
petroleum.  The  relatively  involatile  hydrocarbons  of  which  the 
oil  consists  are  thus  decomposed  ("  cracked  "),  and  gaseous  sub- 
stances of  high  illuminating  power  are  produced.  The  following 
table  shows  the  composition  of  each  of  these  kinds  of  gas,  together 
with  that  of  oil-gas  (Pintsch's),  which  is  composed  entirely  of  the 
products  from  "  cracking  "  oil: 


FIG.  6i. 


Components. 

Coal-Gas. 

Water-Gas. 

Oil-Gas. 

Illuiniiuints             .... 

5.0 

16.6 

45.0 

Heating  gases: 

34.5 

19.8 

38.8 

49.0 

32.1 

14.6 

Carbon  monoxide     .    . 
Impurities: 
Nitrogen  

7.2 
3.2 

26.1 
2.4 

1.1 

Carbon  dioxide      .    .    . 

1.1 

3.0 

Candle  power     

17.5 

25.0 

65.0 

FLAME   AND  ILLUMINANTS  341 

These  are  average  numbers,  and  considerable  variations  from  these 
proportions  are  often  met  with.  The  illuminants  are  unsaturated 
hydrocarbons,  such  as  ethylene  and  acetylene,  and  the  value  of  the 
gas  for  illuminating  purposes  depends  on  the  amount  of  these 
particular  components. 

The    Causes  of  Luminosity   and    Non- Luminosity.  —  The 

study  of  the  chemical  changes  taking  place  in  the  Bunsen  flame, 
particularly  with  the  object  of  explaining  (1)  the  luminosity  of  the 
flame  of  the  pure  gas,  and  (2)  the  non-luminosity  of  that  produced 
by  the  same  gas  when  it  is  mixed  with  air,  has  been  the  subject  of 
many  elaborate  investigations.  The  questions  are:  (1)  Why  is  car- 
bon liberated  in  the  former  case,  and  (2)  why  is  it  not  liberated  in 
the  latter?  Let  us  consider  these  questions  in  order. 

1.  The  investigations  referred  to  show  conclusively  that  the  free 
carbon  in  the  luminous  zone  of  the  ordinary  flame  is  accompanied 
by  free  hydrogen,  and  that  both  are  formed  by  dissociation  of  the 
ethylene.     This  substance,  when  heated,  gives  acetylene,  and  the 
latter  then  dissociates  into  carbon  and  hydrogen  (p.  330) : 

C2H4  ->  H2  +  C2H2  ->  2C  +  H2. 

The  carbon  glows,  until,  as  it  drifts  outwards,  it  encounters  the 
oxygen  of  the  air  and  is  burned.  That  carbon  glows  when  heated 
in  the  absence  of  oxygen,  without  being  consumed,  is  a  fact  familiar 
in  the  behavior  of  the  incandescent  electric  lamp,  the  filament  of 
which  is  made  of  carbon. 

The  conception  that  when  hydrocarbons  burn,  they  first  undergo 
dissociation,  and  then  union  with  oxygen,  is  in  harmony  with  what 
we  have  observed  also  in  the  case  of  the  combustion  of  hydrogen 
sulphide,  where  the  presence  of  free  sulphur  and  free  hydrogen  in  the 
interior  of  the  flame  was  demonstrated  (p.  251). 

2.  The  influence  of  the  air  admitted  to  the  Bunsen  burner,  in 
interfering  with  this  dissociation  in  such  a  way  as  to  destroy  all 
luminosity,  is  the  most  difficult  point  to  explain.     The  effect  is  fre- 
quently attributed  to  the  oxygen  which  the  air  contains.     This  view, 
however,  is  seriously  weakened  by  a  consideration  of  the  undoubted 
fact  that  oxygen  is  not  required.     Carbon  dioxide  and  steam  are 
equally  efficient  when  introduced  instead  of  air.     Even  nitrogen, 
which  cannot  possibly  be  suspected  of  furnishing  any  oxygen,  like- 


342  COLLEGE   CHEMISTRY 

wise  destroys  the  luminosity.  Lewes  has  shown  that  0.5  volumes  of 
oxygen  in  1  volume  of  coal-gas  destroy  the  luminosity.  But  2.30 
volumes  of  nitrogen  or  2.27  volumes  of  air  accomplish  the  same 
result.  Thus  the  efficiency  of  air  is  not  much  greater  than  that  of 
nitrogen,  in  spite  of  the  fact  that  one-fifth  of  the  former  is  oxygen. 
It  is  evident  that  the  effect  is  due,  in  part  at  least,  to  the  dilution 
with  a  cold  gas.  This  is  confirmed  by  the  observation  that  a  cold 
platinum  dish  held  in  a  small  luminous  flame  is  similarly  destructive 
of  the  luminosity.  If  the  tube  of  the  Bunsen  burner  is  heated  so 
that  the  mixed  gases  are  considerably  raised  in  temperature  before 
reaching  the  non-luminous  flame,  the  latter  becomes  luminous.  It 
is  probable,  therefore,  that  the  cold  gas  lowers  the  temperature  of 
the  inner  flame,  and  at  the  same  time  the  dilution  diminishes  the 
speed  with  which  the  free  carbon  is  formed  (Lewes).  Even  if  the 
temperature  is  not  reduced  below  that  at  which  dissociation  of 
the  ethylene  can  occur,  yet  the  dilution  and  cooling,  together,  pre- 
vent that  sharp  dissociation  at  this  particular  point  which  is  neces- 
sary for  the  production  of  the  great  excess  of  free  carbon  needed  to 
furnish  the  light. 

Exercises. —  1.  In  what  way  will  calcium  hydroxide  remove 
hydrogen  sulphide  from  coal-gas  (p.  340)? 

2.  Make  a  section  showing  the  shape  of  the  flame  produced  by 
burning  hydrogen  gas  when  the  latter  issues  from  a  circular  opening. 


CHAPTER  XXXI 
SILICON    AND    BORON 

IN  respect  to  chemical  relations  there  is  a  close  resemblance  between 
silicon  and  carbon.  The  former  element  gives  no  monoxide,  how- 
ever, and  is  quadrivalent  in  all  its  compounds.  In  chemical  character 
it  is  strictly  non-metallic. 

Occurrence.  —  Silicon,  unlike  carbon,  is  not  found  in  the  free 
condition.  In  combination  it  is  the  most  plentiful  element  after 
oxygen,  and  constitutes  more  than  one-quarter  of  the  crust  of  the 
earth.  The  oxide  is  silica  or  sand  (SiO2),  and  this  oxide  and  its 
compounds  are  components  of  many  rocks.  In  the  inorganic  world 
silicon  is  the  characteristic  element  to  almost  as  great  an  extent  as 
is  carbon  in  the  organic  realm. 

Preparation  of  Silicon.  —  When  finely  powdered  magnesium 
and  sand  are  mixed,  and  one  part  of  the  mass  is  heated,  a  violent 
action  spreads  rapidly  through  the  whole: 

2Mg  +  Si02  ->  Si  +  2MgO. 

At  the  same  time,  and  especially  if  excess  of  the  metal  is  used,  some 
magnesium  silicide  Mg2Si  is  formed  also.  The  mixture  is  treated 
with  a  dilute  acid  which  decomposes  the  magnesium  oxide  and 
the  silicide,  and  leaves  the  silicon  (amorphous)  undissolved.  When 
amorphous  silicon  is  dissolved  in  molten  zinc,  the  mass,  after  solidi- 
fication, is  found  to  contain  crystalline  silicon.  The  zinc  is  removed 
by  the  action  of  a  dilute  acid,  the  silicon  remaining  unaffected. 

Properties.  —  Amorphous  silicon  is  a  brown  powder.  It  unites 
with  fluorine  at  the  ordinary  temperature,  with  chlorine  at  430°,  with 
bromine  at  500°,  with  oxygen  at  400°,  with  sulphur  at  600°,  with 
nitrogen  at  about  1000°,  and  with  carbon  and  boron  at  tempera- 
tures attainable  only  in  the  electric  furnace.  It  is  slowly  oxi- 
dized by  aqua  regia  to  silicic  acid,  and  is  dissolved  by  a  mixture  of 

343 


344  COLLEGE  CHEMISTRY 

hydrofluoric  acid  and  nitric  acid,  giving  silicon  tetrafluoride.  Crys- 
tallized silicon  consists  of  black  needles  belonging  to  the  hexagonal 
system.  Its  chemical  behavior  is  like  that  of  the  amorphous  variety 
just  described. 

Both  kinds  of  silicon  act  upon  boiling  solutions  of  potassium  or 
sodium  hydroxide  (cf.  p.  67),  the  ortho-  or  metasilicate  being  formed: 
Si  +  2KOH  +  H2O  ->  K2SiO3  +  2H2. 

Silicon  Hydride  SiH4.  —  Silicon  differs  from  carbon  in  giving 
only  two  well-defined  compounds  with  hydrogen.  The  chief  one 
may  be  liberated  as  a  gas  by  the  action  of  hydrochloric  acid  upon 
magnesium  silicide: 

Mg2Si  +  4HC1  -»  2MgCl2  +  SiH4. 

The  action  is  similar  to  that  by  which  hydrogen  sulphide  is  made. 
The  pure  gas  is  easily  inflammable,  and  unites  with  oxygen  to  form 
water  and  silicon  dioxide.  When  heated  alone,  it  decomposes  into 
its  constituents. 

Carbide  of  Silicon  SiC.  —  This  compound  is  manufactured 
for  use  as  an  abrasive,  and  is  sold  under  the  name  of  carborundum. 
A  mixture  of  quartz-sand,  coke,  and  common  salt  is  heated  to  about 
3500°  in  the  electric  furnace: 

Si02  +  3C  -»  SiC  +  2CO. 

Silicon  carbide  when  pure  is  composed  of  transparent,  colorless, 
hexagonal  plates.  Ordinarily  the  crystals  are  brown  or  black.  It 
stands  next  to  the  diamond  and  carbide  of  boron  in  hardness.  It 
is  used  in  making  machinery  for  polishing  hard  rock,  such  as  granite, 
and  is  employed  also  for  protecting  the  walls  of  puddling  furnaces 

(q.v.),  and  in  other  ways  in  the  steel  industry. 

41 

Silicon  Tetrachloride  and  Tetrafluoride.  —  The  tetrachloride 
SiCl4  is  formed  by  direct  union  of  the  free  elements.  It  is  more 
conveniently  prepared  by  passing  chlorine  over  a  strongly  heated 
mixture  of  silicon  dioxide  and  carbon.  The  gaseous  products  enter 
a  condenser  in  which  the  tetrachloride  assumes  the  liquid  form: 

2C12  +  Si02  +  2C  -»  SiCl4  +  2CO. 

Chlorine  is  unable  to  displace  oxygen  from  combination  with  silicon, 
and  has,  therefore,  when  alone,  no  effect  upon  sand.  In  the  above 


SILICON  345 

action,  therefore,  the-  carbon  is  used  to  secure  the  oxygen  while  the 
chlorine  combines  with  the  silicon.  This  kind  of  interaction  is  in 
some  degree  different  from  any  which  we  have  hitherto  encountered. 
It  bears  no  special  name,  but  the  principle  underlying  it  is  very  com- 
monly employed  (see,  e.g.,  Chlorides  of  boron  and  aluminium). 

Silicon  tetrachloride  is  a  colorless  liquid  (b.-p.  59°)  which  fumes 
strongly  in  moist  air,  and  acts  violently  upon  cold  water,  giving 
silicic  acid: 

SiCl4  +  4H2O  -»  4HC1  +  Si  (OH)  J. 

When  strong  hydrofluoric  acid  acts  upon  sand,  silicon  tetrafluoride 
SiF4  is  liberated: 

Si02  +  4HF  -»  2H2O  +  SiF4. 

Since  the  water  interacts  with  the  tetrafluoride  (see  below),  the 
latter  is  usually  made  by  heating  sand  with  powdered  calcium 
fluoride  and  excess  of  sulphuric  acid.  In  this  way  the  hydrogen 
fluoride  is  generated  in  contact  with  the  sand,  and  at  the  same  time 
the  sulphuric  acid  renders  the  water  inactive.  Hydrofluoric  acid 
acts  in  a  corresponding  way  upon  all  silicates  (q.v.),  whether  these 
are  minerals  or  are  artificial  silicates  like  glass  (cf.  p.  171). 

Silicon  tetrafluoride  is  a  colorless  gas.  It  fumes  strongly  in  moist 
air,  and  acts  vigorously  upon  water.  This  interaction  is  different 
from  that  of  the  tetrachloride,  because  the  excess  of  the  tetrafluoride 
forms  a  complex  compound  with  the  hydrofluoric  acid: 

SiF4  +  4H20  ->  Si(OH)4  (+  4HF)  (1) 

(4HF)  +.  2SiF4  -»  2H2SiF6 (2) 

3SiF4  +  4H2O  ->  Si(OH)4  +  2H2SiF6 

The  silicic  acid  is  precipitated  in  the  water,  and  may  be  separated  by 
filtration,  leaving  a  solution  of  hydrofluosilicic  acid. 

Hydroflwosilicic  Acid  H2SiFG.  —  This  acid  is  stable  only  in 
solution.  When  the  water  is  removed  by  evaporation,  silicon 
tetrafluoride  is  given  off,  while  most  of  the  hydrogen  fluoride  remains 
to  the  last.  Its  salts  are  decomposed  in  a  corresponding  way  when 
they  are  heated.  This  acid  is  used  in  analysis  chiefly  because  its 
potassium  and  sodium  salts  are  amongst  the  few  salts  of  these  metals 
which  are  relatively  insoluble  in  water.  The  barium  salt  is  also 
insoluble,  but  most  of  the  salts  of  the  heavy  metals  are  soluble. 


346  COLLEGE   CHEMISTRY 

Silicon  Dioxide  SiO2.  —  This  substance  may  be  made  in  the 
form  of  a  fine  white  powder  by  heating  precipitated  silicic  acid.  It 
is  found  in  many  different  forms  in  nature.  In  large,  transparent, 
six-sided  prisms  with  pyramidal  ends  it  is  known  as  quartz  or  rock 
crystal.  When  colored  by  manganese  and  iron  it  is  called  amethyst, 
when  by  organic  matter,  smoky  quartz.  A  special  arrangement  of 
the  structure  gives  cat's  eye.  Amorphous  forms  of  the  same  material, 
often  colored  brown  or  red  with  ferric  oxide,  are  agate,  jasper,  and 
onyx,  the  last  much  used  in  making  cameos.  Slightly  hydrated 
forms  of  silica  are  the  opal  and  flint. 

Silicon  dioxide,  although  differing  profoundly  from  carbon  dioxide 
in  its  physical  nature,  nevertheless  behaves  like  the  latter  chemically. 
Thus,  when  boiled  with  potassium  hydroxide  solution  it  forms 
potassium  meta-  or  orthosilicate  : 

Si02  +  4KOH  ->  K4SiO4  +  2H2O. 

The  salt  is  left  as  a  gelatinous  solid  ("  soluble  glass  ")  when  the  water 
is  evaporated.  The  silicates  of  potassium  and  sodium  may  also  be 
obtained  by  boiling  sand  with  the  carbonates  of  these  metals,  carbon 
dioxide  being  displaced.  They  are  produced  more  rapidly,  however, 
usually  as  metasilicates  (see  below),  by  fusing  the  sand  with  the 
alkali  carbonates: 


Si02  +  K2CO3  -»  K2SiO3  +  C0 


2. 


Silica  is  found  in  the  hard  parts  of  straw,  of  some  species  of  horse- 
tail (equisetum),  and  of  bamboo.  In  the  form  of  whetstones  it  is 
used  for  grinding.  The  clear  crystals  are  employed  in  making 
spectacles  and  optical  instruments.  Pure  sand  is  used  in  glass  manu- 
facture (q.v.).  Recently,  small  pieces  of  chemical  apparatus  have 
been  manufactured  by  fusing  quartz  in  the  oxy  hydrogen  flame  or 
the  electric  furnace.  Owing  to  the  low  coefficient  of  expansion  of 
silica,  the  vessels  fashioned  out  of  it  can  be  heated  or  cooled  as 
suddenly  as  we  choose,  without  risk  of  fracture. 

Silicic  Acid  JET4SiO4  —  When  acids  are  added  to  a  solution  of 
sodium  silicate,  silicic  acid  is  set  free.  After  a  little  delay  it  usually 
appears  as  a  gelatinous  precinitate.  When,  however,  the  silicate  is 
poured  into  excess  of  hydrochloric  acid,  no  precipitation  occurs.  The 


SILICON  347 

silicic  acid  remains  in  colloidal  solution  (a  form  of  very  fine  suspen- 
sion).    The  acid  before  precipitation  is  probably  orthosilicic  acid: 

Na4SiO4  +  4HCl-»4NaCl  +  Si(OH)4, 

2HC1  +  H20  ->  2NaCl  +  Si(OH)4, 


but  the  gelatinous  precipitate,  when  dried,  contains  a  smaller  pro- 
portion of  the  elements  of  water.  There  seem  to  be  no  definite 
stages,  indicating  the  existence  of  various  acids,  such  as  we  observe 
with  phosphoric  acid.  The  final  product  of  drying  is  the 
dioxide. 

Silicic  acid  is  a  very  feeble  acid,  and,  therefore,  gives  no  salt  with 
ammonium  hydroxide.  The  potassium  and  sodium  salts  give 
strongly  alkaline  solutions  (cf.  pp.  254,  324). 

Silicates.  —  While  silicic  acid  is  presumed  to  be  the  ortho-acid 
Si  (OH)  4,  and  no  other  silicic  acids  have  been  made,  the  salts  are 
most  easily  classified  by  imagining  them  to  be  derived  from  various 
acids  representing  different  degrees  of  hydration  of  the  dioxide 
(cf.  p.  310),  or,  to  put  it  the  other  way,  different  degrees  of  dehydra- 
tion of  the  ortho-acid.  The  following  equations  show  the  relation 
of  the  ortho-acid  to  some  of  the  silicic  acids  whose  salts  are  most 
commonly  found  amongst  minerals: 

H4SiO4  -    H2O->H2SiO3  (=     H20,Si02)       Metasilicic  acid. 
2H4Si04  -     H20-»HCS207    (-  3H20,2Si02)  \  Disilicic  acids 
2H4Si04-3H20->H2Si205(  =    H2O,2SiO2)  I  J 
3H4SiO4  -  4H2O  -»  H4Si3O8  (=  2H2O,3SiO2)      Trisilicic  acid. 

Di-  and  trisilicates  are  those  derived  from  acids  containing  two  and 
three  units  of  silicic  anhydride,  respectively,  in  the  formula.  The 
valences  of  the  radicals  of  the  acids  are  shown  by  the  number  of 
hydrogen  units  in  the  formulae. 

The  composition  of  minerals  is  often  exceedingly  complex.  This 
is  due  to  the  fact  that  amongst  them  mixed  salts  (p.  231)  are  very 
common,  in  which  the  hydrogen  of  the  imaginary  acid  is  displaced 
by  two  or  more  metals  in  such  a  way  that  the  total  quantity  of  the 
metals  is  equivalent  to  the  hydrogen.  The  following  list  presents 
in  tabular  form  some  typical  or  common  minerals  arranged  according 
to  the  foregoing  classification: 


348  COLLEGE   CHEMISTRY 

(-Zircon,  ZrSi04 
Orthosilicates  (H4SiO4)  ^Mica,  KH2Al3(SiO4)3 

IKaolin,  H2Al2(SiO4)2,H20 

rWollastonite,  CaSiO3 
Metasilicates  (H2SiO3)    j  Beryl,  Gl3Al2(SiO3)6 

iTalc,  H2Mg3(Si03)4 

Disilicate  (H6S2O7)  Serpentine,  Mg3Si2O7, 

Trisilicate  (H4Si3O8)          Orthoclase  (feldspar),  KAlSi3O8 

It  will  be  seen  that  the  total  valence  of  the  metal  units  is  equal  to 
that  of  the  acid  radicals.  Thus,  in  beryl  there  are  six  equivalents  of 
glucinum  (beryllium)  and  six  of  aluminium,  taking  the  place  of 
twelve  units  of  hydrogen  in  (H2Si03)6. 

Mica,  which  is  obtained  in  large  sheets  from  Farther  India,  is  used 
in  making  lamp-chimneys  and  as  an  insulator  in  electrical  apparatus. 
Kaolin,  or  clay,  like  mica,  is  an  acid  orthosilicate. 

Some  of  these  minerals  frequently  occur  mixed  together  as  regular 
components  of  certain  igneous  rocks.  Thus,  granite  is  a  more  or  less 
coarse  mixture  of  quartz,  mica,  and  feldspar.  Frequently  the  oblong, 
flesh-colored  or  white  crystals  of  the  last  are  large  and  very  conspicu- 
ous. Sandstone  is  composed  of  sand  cemented  together  by  clay  or 
by  lime,  and  colored  brown  or  yellow  by  ferric  oxide. 

BORON  B. 

As  regards  chemical  relations,  boron,  being  a  uniformly  trivalent 
element,  is  a  member  of  the  aluminium  family  (see  Table  of  periodic 
system,  at  the  end  of  this  book).  Yet  it  is  a  pronounced  non-metal, 
and  its  oxide  and  hydroxide  are  acidic:  aluminium  is  a  metal,  and 
with  its  oxide  and  hydroxide  basic  properties  predominate.  Boron 
and  its  compounds  really  resemble  carbon  and  silicon  and  their 
compounds  in  all  chemical  properties,  except  the  property  of  valence. 

Occurrence.  —  Like  silicon,  boron  is  found  in  oxygen  compounds, 
namely,  in  boric  acid  (q.v.)  and  its  salts.  Of  the  latter,  sodium  tetra- 
borate  Na^Oj,  or  borax,  came  first  from  India  under  the  name  of 
tincal.  In  constitutes  a  large  deposit  in  Borax  Lake  in  California. 
Colemanite,  Ca2B6On,5H2O,  from  California,  and  other  complex 
borates,  furnish  a  large  part  of  the  commercial  supply  of  compounds 
of  boron. 


BORON  349 

Preparation.  —  When  boric  oxide  is  heated  with  powdered  mag- 
nesium (B203  +  3Mg— >3MgO  +  2B),  black,  amorphous  boron  can 
be  separated  with  some  difficulty  from  the  borides  of  magnesium  in 
the  resulting  mixture.  When  excess  of  powdered  aluminium  is  used, 
hard  crystals  of  boron  are  formed. 

Properties.  —  Boron  unites  with  the  same  elements  as  does 
silicon  (p.  343),  but  with  somewhat  greater  activity.  Like  carbon 
(pp.  257,  299),  it  is  also  oxidized  by  hot,  concentrated  sulphuric  or 
nitric  acid,  the  product  being  boric  acid.  It  interacts  with  fused 
potassium  hydroxide,  giving  a  borate: 

2B  +  6KOH  -*  2K3BO3  +  3H2. 

Boron,  when  heated  with  nitrogen,  unites  directly  to  form  the 
nitride  BN,  a  white  solid.  When  heated  in  the  electric  furnace  with 
carbon,  it  forms  a  carbide  B6C.  This  substance  is  harder  than  car- 
borundum, and  stands  next  to  the  diamond  in  respect  to  hardness. 

The  Halides  of  Boron.  —  By  combined  action  of  carbon  and 
chlorine  on  boric  oxide,  using  the  principle  employed  in  preparing 
silicon  tetrachloride  (p.  344),  the  trichloride  of  boron  BC13  may  be 
made.  It  is  a  liquid  (b.-p.  18°)  which  fumes  strongly  in  moist  air, 
and  is  completely  hydrolyzed  by  water. 

Boron  trifluoride  BF3  is  made  by  the  interaction  of  calcium  fluoride 
and  sulphuric  acid  with  boron  trioxide.  The  mode  of  preparation 
and  the  properties  of  the  substance  recall  silicon  tetrafluoride  (p.  345). 
It  interacts  with  water,  like  the  latter,  giving  boric  acid  and  hydro- 
fluoboric  acid  HBF4: 

4BF3  +  3H20  ->  B(OH)3  +  3HBF4. 

Boric  Acid  and  Boron  Trioxide.  —  Boric  acid  (boracic  acid, 
H3BO3)  is  somewhat  volatile  with  steam,  and  is  found  in  Tuscany 
in  jets  of  water  vapor  (soffioni)  which  issue  from  the  ground. 
Water,  retained  in  basins  of  brickwork,  is  placed  over  the  openings, 
and  from  this  water,  after  evaporation,  boric  acid  is  obtained  in 
crystalline  form.  As  boric  acid  is  very  feeble,  and  withal  little 
soluble,  it  may  also  be  made  by  interaction  of  sulphuric  acid  and 
concentrated  borax  solution: 

Na2B407  +  H2S04  +  5H20  fc;  Na^SO,  +  4H3BO3f. 

Boric  acid  crystallizes  from  water  in  thin  white  plates,  which  are 
unctuous  (like  graphite  and  talc)  to  the  touch.  Its  solubility  in  water 


350  COLLEGE    CHEMISTRY 

is  4  parts  in  100  at  19°,  and  34  in  100  at  100°.  The  solution  scarcely 
affects  litmus.  It  confers  a  green  tint  on  the  Bunsen  flame.  This 
behavior  is  used  as  a  test  for  the  acid.  At  100°  the  acid  slowly  loses 
water,  leaving  metaboric  acid  HB02,  and  at  140°  tetraboric  acid  is 
formed:  4HB02  —  H20  — »  H2B407.  Strong  heating  gives  the 
trioxide  B2O3  a  glassy,  white  solid.  When  dissolved  in  water,  all 
these  dehydrated  compounds  revert  to  boric  acid.  The  solution  of 
boric  acid  in  water  is  used  as  an  antiseptic  in  medicine,  and  as  a 
preservative  for  milk  and  other  foods. 

Borates.  — :  Borates  derived  from  orthoboric  acid  are  practically 
unknown.  The  most  familiar  salt  is  borax  or  sodium  tetraborate. 
The  decahydrate  Na^O^lOH-jO,  which  crystallizes  from  water  at 
27°  in  large,  transparent  prisms,  and  the  pentahydrate  which 
crystallizes  at  56°,  are  both  marketed.  They  are  made  by  crystal- 
lization of  native  borax.  In  Germany,  borax  is  prepared  from 
boracite,  found  at  Stassfurt,  by  decomposing  a  solution  of  the 
mineral  with  hydrochloric  acid: 

MgCl2,2Mg3B8O15  +  12HC1  +  18H20  ^  7MgCl2  +  16B(OH)3. 

The  boric  acid  is  redissolved  in  boiling  water,  and  sodium  carbon- 
ate is  added:  4B(OH)3  +  Na2CO3  ->  Na2B407  +  6H20  +  CO2.  In 
California  it  is  made  from  colemanite  by  interaction  with  sodium 
carbonate. 

Since  boric  acid  is  a  feeble  acid,  borax  is  extensively  hydrolyzed  by 
water,  and  the  solution  has  a  marked  alkaline  reaction  (cf.  p.  254). 

When  heated  with  oxides  of  metals,  sodium  tetraborate  behaves 
like  sodium  metaphosphate  (cf.  p.  312),  and  is  used  in  the  form  of 
beads  in  analysis.  If  its  formula  be  written  2NaBO2,B203  (cf.  p. 
310)  it  will  be  seen  that  a  considerable  excess  of  the  acid  anhydride 
is  contained  in  it,  and  that,  therefore,  a  mixed  metaborate  may  be 
formed  by  union  with  some  basic  oxide.  Thus,  with  a  trace  of 
cupric  oxide,  the  bead  is  tinged  with  blue,  from  the  presence  of  a 
compound  like  2NaB02,Cu(B02)2.  Cobalt  compounds  give  a  deep- 
blue  color  to  the  bead. 

Exercises. —  1.  Compare  and  contrast  the  elements  carbon  and 
silicon,  and  their  corresponding  compounds. 

2.  What  would  be  the  interaction  between  aqueous  solutions  of 
an  ammonium  salt  and  of  sodium  orthosilicate  (cf.  p.  347)? 


CHAPTER  XXXII 
THE   BASE-FORMING    ELEMENTS 

IN  the  present  chapter  a  preliminary  view  of  the  chemistry  of  the 
metallic  elements  is  given. 

Physical  Properties  of  the  Metals.  —  Metals  show  what  is 
commonly  called  a  metallic  luster,  but,  as  a  rule,  they  do  so  only 
when  in  compact  form.  Magnesium  and  aluminium  exnibit  it 
when  powdered,  but  most  of  the  metals  when  in  this  condition  are 
black. 

The  metals  can  all  be  obtained  in  crystallized  form,  when  a  fused 
mass  is  allowed  to  cool  slowly  and  the  unsolidified  portion  is  poured 
off.  In  almost  all  cases  the  crystals  belong  to  the  regular  system. 

The  metals  vary  in  specific  gravity  from  lithium,  which  is  little 
more  than  half  as  heavy  as  water  (sp.  gr.  0.59),  to  osmium,  whose 
specific  gravity  is  22.5.  Those  which  have  a  specific  gravity  less 
than  5,  namely,  potassium,  sodium,  calcium,  magnesium,  aluminium, 
and  barium,  are  called  the  light  metals,  and  the  others  the  heavy 
metals. 

Most  metals  are  malleable,  and  can  be  beaten  into  thin  sheets  with- 
out loss  of  continuity.  Those  which  are  allied  to  the  non-metals, 
however,  such  as  arsenic,  antimony,  and  bismuth,  are  brittle.  The 
order  of  the  elements  in  respect  to  this  property,  beginning  with  the 
most  malleable,  is:  Au,  Ag,  Cu,  Sn,  Pt,  Pb,  Zn,  Fe,  Ni. 

The  tenacity  of  the  metals  places  them  in  an  order  different  from 
the  above.  It  is  measured  by  the  number  of  kilograms  which  a 
piece  of  the  metal  1  sq.  mm.  in  section  can  sustain  without  breaking. 
The  values  are  as  follows:  Fe  62,  Cu  42,  Pt  34,  Ag  29,  Au  27,  Al  20, 
Zn  5,  Pb  2. 

The  hardness  is  measured  by  the  ease  with  which  the  material  may 
be  disintegrated  by  a  sharp,  hard  instrument.  Potassium  is  as  soft 
as  wax,  while  chromium  is  hard  enough  to  cut  glass. 

The  temperature  at  which  the  metal  fuses  has  an  important  bearing 

351 


352 


COLLEGE   CHEMISTRY 


on  its  manufacture.    Most  of  the  following  melting-points  are  only 
approximate: 


Mercury  . 

-40° 

Zinc     .    .   . 

420° 

Cast  iron 

1150° 

Potassium 

62° 

Antimony 

437° 

Nickel  .    . 

1500° 

Sodium     . 

96° 

Aluminium 

700° 

Iron  (pure) 

1550° 

Tin    ... 

230° 

Magnesium 

750° 

Platinum  . 

1780° 

Bismuth  . 

264° 

Silver      .    . 

954° 

Manganese 

1900° 

Cadmium 

320° 

Copper   .    . 

1057° 

Chromium 

2000° 

Lead     .   . 

326° 

Gold    .    .    . 

1075° 

Iridium    . 

2200° 

It  will  be  seen  that  mercury  is  a  liquid,  that  potassium  and  sodium 
melt  below  the  boiling-point  of  water,  and  that  the  metals  down  to 
the  foot  of  the  second  column  can  be  melted  easily  with  the  Bunsen 
flame. 

The  methods  of  manufacture  and  the  treatment  of  metals  are 
much  influenced  also  by  their  volatility.  The  following  are  easily 
distilled:  Mercury,  b.-p.  357°;  potassium  and  sodium,  b.-p.  about 
700°;  cadmium,  b.-p.  770°;  zinc,  b.-p.  950°.  Even  the  most  invola- 
tile  metals  can  be  converted  into  vapor  in  the  electric  arc. 

In  many  cases  molten  metals  dissolve  in  one  another  freely.  The 
results  are  called  alloys,  and  in  some  cases  have  the  properties  of 
solid  solutions.  Sometimes,  as  in  the  case  of  lead  and  tin,  mixtures 
can  be  formed  in  all  proportions.  On  the  other  hand,  the  solubility 
may  be  limited,  as  in  the  case  of  zinc  and  lead,  where  only  1.6  parts 
of  the  former  dissolve  in  100  parts  of  the  latter.  The  colors  of  alloys 
are  not  the  average  of  those  of  the  constituents.  Thus,  the  nickel 
alloy  used  in  coining  contains  75  per  cent  of  copper  and  25  per  cent 
of  nickel,  yet  it  shows  none  of  the  color  of  the  former. 

Alloys  in  which  mercury  forms  one  of  the  components  are  known 
as  amalgams  (Gk.  /AoAay/xa,  a  soft  mass),  and  are  formed  with  especial 
ease  by  the  lighter  metals.  Of  the  common  metals,  iron  is  the  least 
miscible  with  mercury. 

The  good  conductivity  of  metals  for  electricity  distinguishes  them 
with  some  degree  of  sharpness  from  the  non-metals.  They  show 
considerable  variation  amongst  themselves,  silver  conducting  sixty 
times  as  well  as  mercury.  The  following  table  gives  the  conduc- 
tivities of  the  metals  (expressed  in  terms  of  the  number  of  meters  of 
wire  1  sq.  mm.  in  section  which,  at  15°,  offer  a  resistance  of  one 
ohm): 


THE   BASE-FORMING   ELEMENTS  353 

Silver,  cast 62.89  Nickel,  cast 7.59 

Copper,  commercial  .    .57.40  Iron,  drawn 7.55 

Gold,  cast 46.30  Platinum 5.7-8.4 

Aluminium,  commercial  31.52  Steel 5.43 

Zinc,  rolled 16.95  Lead      4.56 

Brass 14.17  Mercury 1.049 

To  compare  these  conductivities  with  those  of  solutions,  it  may 
be  said  that  decinormal  hydrochloric  acid  (p.  226)  has  a  conductivity 
on  the  above  scale  of  0.035,  or  a  thirtieth  of  that  of  mercury. 

General  Chemical  Relations  of  the  Metallic   Elements.  — 

Since  most  of  the  compounds  of  the  metals  are  ionogens,  their  solu- 
tions, except  when  the  metal  is  a  part  of  a  compound  ion,  all  contain 
the  metal  in  the  ionic  state,  and  the  resulting  substances,  such  as 
potassium-ion  and  cupric-ion,  have  constant  properties,  irrespective 
of  the  nature  of  the  negative  ion  with  which  they  may  be  mixed. 
The  properties  of  the  ions,  simple  and  compound,  are  much  used  in 
making  tests  in  analytical  chemistry.  On  the  other  hand,  the 
chemical  properties  of  the  oxides  and  of  the  salts  in  the  dry  state  are 
of  importance  in  connection  with  metallurgy. 

There  are  three  chemical  properties  which  are  characteristic  of  the 
metallic  elements.  The  first  two  of  them  have  already  been  discussed 
somewhat  fully. 

1.  The  metals  are  able  by  themselves  to  form  positive  radicals  of 
salts,  and,  therefore,  to  exist  alone  as  positive  ions  (pp.  224,  271). 

2.  The  oxides  and  hydroxides  of  the  metals  are  basic  (pp.  81,  271). 

3.  Each  typical  metal  has  at  least  one  halogen  compound  which 
is  little,  if  at  all,  hydrolyzed  by  water  (p.  272).     The  same  thing  is 
true  of  nitrates  and  other  salts  of  active  acids. 

In  reference  to  the  third  characteristic,  the  non-hydrolysis  of  halides 
of  typical  metals,  a  word  of  explanation  is  required.  Active  bases 
(hydroxides  of  typical  metals),  such  as  sodium  hydroxide,  give,  with 
feeble  acids,  such  as  H2S  (p.  253),  H2PO/  (p.  311),  H2CO3  (p.  324), 
H2SiO3  (p.  347)  and  H3BO3  (p.  350),  salts  whose  solutions  are  alkaline 
in  reaction.  This  is  due  to  hydrolysis.  But  active  bases  give  with 
active  acids,  such  as  HC1,  and  HNO3,  salts  whose  solutions  are  neutral 
in  reaction.  This  is  the  fact  expressed  in  the  third  characteristic  of 
the  metallic  elements.  The  less  active  bases,  being  hydroxides  of 
less  active  metallic  elements,  give,  with  active  acids,  salts  whose 
solutions  are  not  neutral,  but  acid  in  reaction.  Thus  cupric  chloride 


354  COLLEGE   CHEMISTRY 

solution  is  feebly  acid.  This  is  because  there  is  a  tendency  for  the 
ions  of  the  water  to  form  the  slightly  dissociated  molecules  of  the 

Cu"  +  20H'  +  2H'  -»  Cu(OH)2  +  2H*. 

Finally,  a  salt  derived  from  a  base  and  an  acid,  both  of  which  are 
weak  is  also  hydrolyzed,  often  completely.  Aluminium  carbonate 
and  ammonium  silicate  (p.  347)  are  examples  of  salts  which,  for  this 
reason,  are  completely  hydrolyzed.  The  resulting  mixture  may  have 
an  acid  or  a  basic  reaction,  if  the  acid  or  the  base  is  sufficiently 
soluble  and  sufficiently  active.  Thus,  ammonium  sulphide  solution 
is  alkaline. 

Aside  from  these  points,  many  features  in  the  behavior  of  metals 
and  their  compounds  are  summed  up  in  the  electromotive  series 
(p.  245).  The  reader  should  re-read  all  the  parts  referred  to  above 
before  proceeding  farther.  He  should  also  reexamine  the  various 
kinds  of  chemical  changes  discussed  on  pp.  124,  163,  189  and  par- 
ticularly the  varieties  of  ionic  chemical  change  on  p.  243. 

Occurrence  of  the  Metals  in  Nature.  —  The  minerals  from 
which  metals  are  extracted  are  known  as  ores.  They  present  a  com- 
paratively small  number  of  different  kinds  of  compounds.  Most  of 
the  metals  are  found  in  more  than  one  of  these  forms,  so  that  in  the 
following  statement  the  same  metal  frequently  occurs  more  than 
once. 

When  the  metal  occurs  free  in  nature  it  is  said  to  be  native.  Thus 
we  have  gold,  silver,  metals  of  the  platinum  group,  copper,  mercury, 
bismuth,  antimony,  and  arsenic  occurring  native  (cf.  p.  245). 

The  metals  whose  oxides  are  important  minerals  are  iron,  man- 
ganese, tin,  zinc,  copper,  and  aluminium.  The  metals  are  obtained 
commercially  from  the  oxides  in  each  of  these  cases. 

The  metals  whose  sulphides  are  used  as  ores  are  nickel,  cobalt, 
antimony,  lead,  cadmium,  zinc,  and  copper. 

From  the  carbonates  we  obtain  iron,  lead,  zinc,  and  copper. 
Several  other  metals,  such  as  manganese,  magnesium,  barium, 
strontium,  and  calcium  occur  in  larger  or  smaller  quantities  in  the 
same  form  of  combination. 

The  metals  which  occur  as  sulphates  are  those  whose  sulphates  are 
not  freely  soluble,  namely,  lead,  barium,  strontium,  and  calcium. 

Compounds  of  metals  with  the  halogens  are  not  so  numerous. 


THE  BASE-FORMING  ELEMENTS  355 

Silver  chloride  furnishes  a  limited  amount  of  silver.  Sodium  and 
potassium  chlorides  are  found  in  the  salt-beds,  and  cryolite  3NaF, 
A1F3  is  used  in  the  manufacture  of  aluminium. 

The  natural  silicates  are  very  numerous,  but  are  seldom  used  for 
the  preparation  of  the  metals.  Many  of  them  are  employed  for 
other  commercial  purposes,  kaolin  (p.  348)  being  a  conspicuous 
example  of  this  class. 

Methods  of  Extraction  from  the  Ores.  —  The  art  of  extract- 
ing metals  from  their  ores  is  called  metallurgy.  Where  the  metal  is 
native,  the  process  is  simple,  since  melting  away  from  the  matrix 
(p.  248)  is  all  that  is  required.  Frequently  a  flux  is  added.  A  flux 
usually  is  a  substance  which  interacts  with  infusible  materials  to 
give  fusible  ones.  It  combines  with  the  matrix,  giving  a  fusible 
slag.  Since  the  slag  is  a  melted  salt,  usually  a  silicate,  and  does  not 
mix  at  all  with  the  molten  metal,  separation  of  the  products  is 
easily  effected.  When  the  ore  is  a  compound,  the  metal  has  to  be 
liberated  by  our  furnishing  a  material  capable  of  combining  with  the 
other  constituent.  The  details  of  the  process  depend  on  various 
circumstances.  Thus  the  volatile  metals,  like  zinc  and  mercury,  are 
driven  off  in  the  form  of  vapor,  and  secured  by  condensation.  The 
involatile  metals,  like  copper  and  iron,  run  to  the  bottom  of  the 
furnace  and  are  tapped  off. 

Where  the  ore  is  an  oxide  it  is  usually  reduced  by  heating  with 
carbon  in  some  form.  This  holds  for  the  oxides  of  iron  and  copper, 
for  example.  Some  oxides  are  not  reducible  by  carbon  in  an  ordinary 
furnace.  Such  are  the  oxides  of  calcium,  strontium,  barium,  mag- 
nesium, aluminium,  and  the  members  of  the  chromium  group.  At 
the  temperature  of  the  electric  furnace  even  these  may  be  reduced, 
but  the  carbides  are  formed  under  such  circumstances,  and  the 
metals  are  more  easily  obtained  otherwise.  Recently,  heating  the 
pulverized  oxide  with  finely  powdered  aluminium  has  come  into  use, 
particularly  for  operations  on  a  small  scale.  Iron  oxide  is  easily 
reduced  by  this  means,  and  even  the  metals  manganese  and  chro- 
mium may  be  liberated  from  their  oxides  quite  readily  by  this  action. 
This  procedure  has  received  the  name  aluminothermy  (q.v .)  on  account 
of  the  great  amounts  of  heat  liberated.  In  the  laboratory  the  oxides 
of  the  less  active  metals  are  frequently  reduced  in  a  stream  of 
hydrogen  (cf.  p.  244). 


356  COLLEGE    CHEMISTRY 

When  the  ore  is  a  carbonate,  it  is  first  heated  strongly  to  drive  out 
the  carbon  dioxide  (cf.  p.  322) .  FeCO3  <=>  FeO  +  CO2  f,  and  then  the 
oxide  is  treated  according  to  one  of  the  above  mentioned  methods. 
When  the  ore  is  a  sulphide,  it  has  to  be  roasted,  or  calcined  (p.  256), 
in  order  to  remove  the  sulphur,  and  the  resulting  oxide  is  then  treated 
as  described  above. 

Chlorides  and  fluorides  of  the  metals  can  be  decomposed  by  heating 
with  metallic  sodium.  This  method  was  formerly  employed  in  the 
making  of  magnesium  and  aluminium. 

The  metals  which  are  not  readily  secured  in  any  of  the  above  ways, 
can  be  obtained  easily  by  electrolysis  of  the  fused  chloride  or  of  some 
other  simple  compound.  Aluminium  is  now  manufactured  entirely 
by  the  electrolysis  of  a  solution  of  aluminium  oxide  in  molten  cryolite. 

Compounds  of  the  Metals  :  Oxides  and  Hydroxides.  —  The 
oxides  may  be  made  by  direct  burning  of  the  metal,  by  heating  the 
nitrates  (cf.  p.  296),  the  carbonates  (cf.  p.  322),  or  the  hydroxides: 
Ca(OH)2  «=±  CaO  +  H2O  f .  They  are  practically  insoluble  in  water, 
although  the  oxides  of  the  metals  of  the  alkalies  and  of  the  metals  of 
the  alkaline  earths  interact  with  water  rapidly  to  give  the  hydrox- 
ides. Oxides  are  usually  stable.  Those  of  gold,  platinum,  mercury, 
and  silver  decompose  when  heated,  yet  with  increasing  difficulty  in 
this  order.  The  metals,  like  the  non-metals,  frequently  give  several 
different  oxides.  Those  of  the  univalent  metals  (having  the  form 
K2O),  if  we  leave  cuprous  oxide  and  aurous  oxide  out  of  account, 
have  the  most  strongly  basic  qualities.  Those  of  the  bivalent  metals 
of  the  form  MgO,  when  this  is  the  only  oxide  which  they  furnish,  are 
base-forming.  Those  of  the  trivalent  metals  of  the  form  A12O3, 
known  as  sesquioxides  (Lat.  sesqui-,  one-half  more),  are  the  least 
basic  of  the  basic  oxides.  The  oxides  of  the  forms  Sn02,  Sb205, 
CrO3,  and  Mn2O7,  in  which  the  metals  have  valences  from  4  to  7,  are 
mainly  acid-forming  oxides,  although  the  same  elements  usually 
have  other  lower  oxides,  which  are  basic. 

The  hydroxides  are  formed,  in  the  cases  of  the  metals  of  the  alkalies 
and  alkaline  earths,  by  direct  union  of  water  with  the  oxides.  They 
are  produced  also  by  double  decomposition  when  a  soluble  hydroxide 
acts  upon  a  salt  (cf.  p.  239).  All  hydroxides,  except  those  of  the 
alkali  metals,  lose  the  elements  of  water  when  heated,  and  the  oxide 
remains.  In  some  cases  the  loss  takes  place  by  stages,  just  as  was 


THE   BASE-FORMING  ELEMENTS.  357 

the  case  with  orthophosphoric  acid  (p.  309).  Thus  lead  hydroxide 
Pb(OH)2  (q.v.)  first  gives  the  hydroxide Pb2O(OH)2,  then  Pb3O2 (OH) 3, 
and  then  the  oxide  PbO.  The  hydroxides,  with  the  exception  of 
those  of  the  metals  of  the  alkalies  and  alkaline  earths,  are  all  little 
soluble  in  water.  The  hydroxides  of  mercury  and  silver,  if  they  are 
formed  at  all,  are  evidently  unstable,  for,  when  either  material  is 
dried,  it  is  found  to  contain  nothing  but  the  corresponding  oxide. 

Compounds  of  the  Metals :  Salts.  —  It  may  be  said,  in  general, 
that  each  mental  may  form  a  salt  by  combination  with  each  one  of 
the  acid  radicals.  In  the  succeeding  chapters  we  shall  describe  only 
those  salts  which  are  manufactured  commercially,  or  are  of  special 
interest  for  some  other  reason.  The  various  salts  will  be  described 
under  each  metal.  Here,  however,  a  few  remarks  may  be  made 
about  the  characteristics  of  the  more  common  groups  of  salts.  The 
salts  are  classified  according  to  the  acid  radicals  which  they  contain. 

The  chlorides  may  be  made  by  the  direct  union  of  chlorine  with 
the  metal  (cf.  p.  113),  or  by  the  combined  action  of  carbon  and 
chlorine  upon  the  oxide  (cf.  p.  344).  The  latter  method  is  used  in 
making  chromium  chloride.  The  general  methods  for  making  any 
salt  (p.  123),  such  as  the  interaction  of  a  metal  with  an  acid,  or  of  the 
oxide,  hydroxide,  or  another  salt  with  an  acid,  or  the  double  decom- 
position of  two  salts,  may  be  used  also  for  making  chlorides.  The 
chlorides  are  for  the  most  part  soluble  in  water.  Silver  chloride, 
mercurous  chloride,  and  cuprous  chloride  are  almost  insoluble,  how- 
ever, and  lead  chloride  is  not  very  soluble.  Most  of  the  chlorides  of 
metals  dissolve  without  decomposition,  but  hydrolysis  is  conspicuous 
in  the  case  of  the  chlorides  of  the  trivalent  metals,  such  as  aluminium 
chloride  and  ferric  chloride  (cf.  p.  353).  The  chlorides  of  some  of 
the  bivalent  metals  are  hydrolyzed  also,  but,  as  a  rule,  only  when 
they  are  heated  with  water.  This  is  the  case  with  the  chlorides  of 
magnesium,  calcium,  and  zinc.  Most  of  the  chlorides  are  stable 
when  heated,  but  those  of  the  noble  metals,  particularly  gold  and 
platinum,  are  decomposed,  and  chlorine  escapes.  The  chlorides  are 
usually  the  most  volatile  of  the  salts  of  a  given  metal,  and  so  are 
preferred  for  the  production  of  the  spectrum  (q.v.)  of  the  metal,  and 
for  fixing  the  atomic  weight  of  the  metal  by  use  of  the  vapor  density. 
Some  of  the  metals  form  two  or  more  different  chlorides.  For 
example,  indium  gives  InCl,  InCl2,  and  InCl3. 


358  COLLEGE   CHEMISTRY 

The  sulphides  are  formed  by  the  direct  union  of  the  metal  with 
sulphur,  or  by  the  action  of  hydrogen  sulphide  or  of  some  soluble 
sulphide  upon  a  solution  of  a  salt  (cf.  p.  254).  In  one  or  two  cases 
they  are  made  by  the  reduction  of  the  sulphate  with  carbon.  The 
sulphides,  except  those  of  the  alkali  metals,  are  but  little  soluble  in 
water.  The  sulphides  of  aluminium  and  chromium  are  hydrolyzed 
completely  by  water,  giving  the  hydroxides,  and  those  of  the  metals 
of  the  alkaline  earths  are  partially  hydrolyzed  (cf.  p.  254). 

The  carbides  are  usually  formed  in  the  electric  furnace  by  inter- 
action of  an  oxide  with  carbon  (cf.  p.  321).  Some  'of  them  are 
decomposed  by  contact  with  water,  after  the  manner  of  calcium 
carbide,  giving  a  hydroxide  and  a  hydrocarbon.  Of  this  class  are 
lithium  carbide  Li2C2,  barium  and  strontium  carbides  BaC2  and  SrC2, 
aluminium  carbide  A14C3,  manganese  carbide  MnC,  and  the  carbides 
of  potassium  and  glucinum.  Others,  such  as  those  of  molybdenum 
CMo2  and  chromium  Cr3C2,  are  not  affected  by  water. 

The  nitrates  may  be  made  by  any  of  the  methods  used  for  prepar- 
ing salts.  They  are  all  at  least  fairly  soluble  in  water. 

The  sulphates  are  made  by  the  methods  used  for  making  salts,  and 
in  some  cases  by  the  oxidation  of  sulphides.  They  are  all  soluble 
in  water,  with  the  exception  of  those  of  lead,  barium,  and  strontium. 
Calcium  sulphate  is  meagerly  soluble. 

The  carbonates  are  prepared  by  the  methods  used  for  making  salts. 
They  are  all  insoluble  in  water,  with  the  exception  of  those  of  sodium 
and  potassium.  The  hydroxides  of  aluminium  and  tin  are  so  feebly 
basic  that  these  metals  do  not  form  stable  carbonates  (cf.  pp.  347, 
354). 

The  phosphates  and  silicates  are  prepared  by  the  methods  used  in 
making  salts.  The  former  are  obtained  also  by  special  processes 
already  described  (p.  312).  With  the  exception  of  the  salts  of 
sodium  and  potassium,  all  the  salts  of  both  these  classes  are  insoluble. 

For  the  exact  solubilities  of  a  large  number  of  bases  and  salts  at  189, 
see  the  Table  inside  the  cover,  at  the  front  of  this  book.  Solubilities 
at  all  temperatures  are  shown  in  the  diagram,  Fig.  40,  p.  104. 

Exercises. —  1.  What  do  we  mean  by  saying  that  an  oxide  is 
strongly  or  feebly  basic,  or  that  it  is  acidic? 

2.  What  is  meant  by  the  same  terms  when  applied  to  an 
hydroxide? 


THE  BASE-FORMING  ELEMENTS  359 

3.  Compare  the  molar  solubilities  at  18°,  (a)  of  the  halides  of 
silver,  and  (b)  of  the  carbonates  and  (c)  oxalates  of  the  metals  of  the 
alkaline  earths,  noting  the  relation  between  solubility  and  atomic 
weight. 

4.  What  is  the  molar  concentration  of  chloride-ion  in  saturated 
solutions  of  silver  chloride  and  lead  chloride  at  18°,  assuming  com- 
plete ionization  in  these  very  dilute  solutions? 


CHAPTER  XXXIII 

THE    METALLIC    ELEMENTS    OF    THE    ALKALIES: 
POTASSIUM    AND    AMMONIUM 

THE  metals  of  this  family,  with  their  atomic  weights,  are : 

Lithium 7.0    Rubidium 85.5 

Sodium       23.0    Caesium 132.9 

Potassium      39 . 1 

The  Chemical  Relations  of  the  Metallic  Elements  of  the 
Alkalies.  —  The  metals  which  are  chemically  most  active  are 
included  in  this  group,  and  the  activity  increases  with  rising  atomic 
weight,  caesium  being  the  most  active  positive  element  of  all.  A 
freshly  cut  surface  of  any  of  these  metals  tarnishes  by  oxidation  as 
soon  as  it  is  exposed  to  the  air.  All  of  these  metals  decompose 
water  violently  (cf.  p.  66),  liberating  hydrogen.  The  hydroxides 
which  are  formed  by  this  action  are  exceedingly  active  bases,  that 
is  to  say,  they  give  a  relatively  large  concentration  of  hydroxide-ion 
in  solutions  of  a  given  molecular  concentration  (p.  229).  In  the  dry 
form  these  hydroxides  are  not  decomposed  by  heating,  while  the 
hydroxides  of  all  other  metals  lose  water  more  or  less  easily.  In  all 
their  compounds  the  metals  of  the  alkalies  are  univalent. 

The  compounds  of  ammonium  will  be  discussed  in  connection 
with  those  of  potassium,  to  which  they  present  the  greatest  resem- 
blance. 

The  solubilities  are  often  decisive  factors  in  connection  with  the 
preparation  and  use  of  salts.  The  reader  will  find  most  of  these  in 
the  table  on  the  inside  of  the  cover,  at  the  front  of  this  book,  or  in  the 
diagram  on  p.  104,  and,  as  a  rule,  the  values  will  not  be  repeated  in 
the  descriptive  paragraphs. 

POTASSIUM. 

Occurrence.  —  Silicates  containing  potassium,  such  as  feldspar 
and  mica  (p.  348),  are  constant  constituents  of  volcanic  rocks. 
These  minerals  are  not  used  commercially  as  sources  of  potassium 

360 


POTASSIUM  361 

compounds.  The  salt  deposits  (see  below)  contain  potassium 
chloride,  alone  (sylvite)  and  in  combination  with  other  salts,  and 
most  of  the  compounds  of  potassium  are  manufactured  from  this 
material.  Part  of  our  potassium  nitrate,  however,  is  purified  Bengal 
saltpeter  (p.  292).  Potassium  sulphate  occurs  also  in  the  salt  layers. 

Preparation.  —  Potassium  was  first  made  by  Davy  (1807)  by 
bringing  the  wires  from  a  battery  in  contact  with  a  piece  of  moist 
potassium  hydroxide.  Globules  of  the  metal  appeared  at  the  nega- 
tive wire.  Electrolytic  processes  have  just  come  into  use  commer- 
cially, molten  potassium  chloride  being  the  substance  decomposed. 
Castner's  reduction  process  involves  the  heating  of  potassium  hy- 
droxide with  a  spongy  mass  which  is  essentially  a  carbide  of  iron 
(CFe2) .  The  latter  is  made  by  heating  together  pitch  and  iron  filings  : 

6KOH  +  2C  -+  2K2C03  +  3H2  +  2K. 
The  potassium  passes  off  as  vapor,  and  is  condensed. 

Physical  and  Chemical  Properties.  —  Potassium  is  a  silver- 
white  metal  which  melts  at  62.5°.  It  boils  at  667°,  giving  a  green 
vapor. 

The  density  of  the  vapor  shows  the  molecular  weight  of  potassium 
to  be  about  40,  so  that  the  vapor  is  a  monatomic  gas.  The  element 
unites  violently  with  the  halogens,  sulphur,  and  oxygen.  In  con- 
sequence of  the  latter  fact  it  is  usually  kept  under  petroleum,  an  oil 
which  neither  contains  oxygen  itself,  nor  dissolves  a  sufficient  amount 
of  oxygen  from  the  air  to  permit  much  oxidation  of  the  potassium  to 
take  place.  A  white,  crystalline  hydride  KH  is  formed  when  hy- 
drogen is  passed  over  potassium  heated  to  360°.  When  thrown  into 
water  it  gives  potassium  hydroxide,  and  the  hydrogen  is  liberated. 

Potassium  Chloride  KCl. —  Sea-water  and  the  waters  of  salt 
lakes  contain  a  relatively  small  proportion  of  potassium  compounds. 
During  the  evaporation  of  such  waters,  however,  the  potassium 
compounds  tend  to  accumulate  in  the  mother-liquor  while  sodium 
chloride  is  being  deposited.  Hence  the  upper  layers  of  salt  deposits 
are  the  richest  in  compounds  of  potassium.  Thus,  at  Stassfurt,  near 
Magdeburg,  there  is  a  thickness  of  more  than  a  thousand  meters  of 
common  salt.  Above  this  are  25-30  meters  of  salt  layers  in  which 
the  potassium  salts  are  chiefly  found. 


362  COLLEGE    CHEMISTRY 

The  chief  forms  in  which  potassium  chloride  is  found  in  the  salt 
beds  are  sylvite  (KC1)  and  carnallite  (KCl,MgCl2,6H2O).  The  latter 
salt  is  heated  with  a  small  amount  of  water,  or  with  a  mother- 
liquor  obtained  from  a  previous  operation  and  containing  sodium 
and  magnesium  chlorides.  The  magnesium  sulphate  which  it 
contains  as  an  impurity  remains  undissolved.  From  the  clear 
liquid,  when  it  cools,  potassium  chloride  is  deposited  first  and  then 
carnallite.  The  former  is  taken  out  and  purified,  and  the  latter  goes 
through  the  process  again.  This  potassium  chloride  is  the  source 
from  which  most  of  our  potassium  hydroxide  and  potassium  car- 
bonate, as  well  as  salts  of  minor  commercial  importance,  are  made. 
It  is  a  white  substance  crystallizing  in  cubes,  melting  at  about  750°, 
and  slightly  volatile  at  high  temperatures. 

The  Other  Halides  of  Potassium.  —  When  iodine  is  heated 
in  a  strong  solution  of  potassium  hydroxide,  potassium  iodate  and 
potassium  iodide  are  both  formed  (p.  198) : 

6KOH  +  3I2  ->  5KI  +  KI03  +  3H2O. 

The  dry  residue  from  evaporation  is  heated  with  powdered  carbon  to 
reduce  the  iodate,  and  all  the  iodide  can  then  be  purified  by  recrys- 
tallization.  Another  method  of  preparation  consists  in  rubbing 
together  iodine  and  iron  filings  under  water.  The  soluble  ferrous 
iodide  (FeI2)  thus  formed  is  then  treated  with  additional  iodine  and 
gives  a  substance  Fe3I8,  intermediate  in  composition  between  ferrous 
and  ferric  iodides.  This  is  also  soluble.  When  potassium  car- 
bonate is  added  to  the  solution,  a  hydrated  magnetic  oxide  of  iron 
is  precipitated,  carbon  dioxide  escapes,  and  evaporation  of  the 
filtered  solution  gives  potassium  iodide : 

Fe3I8  4-  4K2CO3  +  4H2O  ->  SKI  +  Fe3(OH)8  +  4CO2. 

The  salt  forms  large,  somewhat  opaque  cubes  (m.-p.  623°).  It  is 
used  in  medicine  and  in  photography  (q.v.). 

Potassium  bromide  KBr  may  be  made  in  either  of  the  ways  used 
for  the  iodide.  It  crystallizes  in  cubes.  It  is  used  in  medicine  and 
for  precipitating  silver  bromide  in  making  photographic  plates  (q.v.). 

The  fluoride  of  potassium  K2F2  may  be  obtained  by  treating  the 
carbonate  or  hydroxide  with  hydrofluoric  acid.  It  is  a  deliquescent, 
white  salt.  When  treated  with  an  equi-molecular  quantity  of  hydro- 
fluoric acid  it  forms  potassium-hydrogen  fluoride  KHF2,  a  white  salt 
which  is  also  very  soluble. 


POTASSIUM 


363 


Potassium  Hydroxide  KOH. —  This  compound,  known  also  as 
caustic  potash,  and  sometimes  as  potassium  hydrate,  was  formerly 
made  entirely  by  boiling  potassium  carbonate  with  calcium  hydrox- 
ide suspended  in  water  (milk  of  lime) : 


Ca(OH)2(solid)  <=»Ca(OH)2(dslvd)  ^ 


Ca" 


:CaCO3(dslvd); 


±CaCO3 

(solid). 

The  operation  is  conducted  in  iron  vessels,  because  porcelain,  being 
composed  of  silicates,  interacts  with  solutions  of  bases.  On  account 
of  the  very  limited  solubility  of  the  calcium  hydroxide  (0.17  g.  in 
100  g.  Aq),  the  water  takes  up  fresh  portions  into  solution  only  when 
the  part  dissolved  has  already  undergone  chemical  change.  The 
calcium  carbonate  which  is  precipitated  is,  however,  still  more 
insoluble  (0.0013  g.  in  100  g.  Aq)  than  is  the  hydroxide,  and  hence 
the  action  goes  forward.  After  the  precipitate  has  settled,  the 
potassium  hydroxide  is  obtained  by  evaporation  of  the  clear  liquid, 
K*  +  OH'  ->  KOH. 

Recently  much  potassium  hydroxide  has  been  manufactured  by 


Fra.  62. 


electrolytic  processes.  When  a  solution  of  potassium  chloride  is 
electrolyzed,  chlorine  is  liberated  at  the  anode,  and  hydrogen  and 
potassium  hydroxide  at  the  cathode.  These  two  sets  of  products 
must  be  kept  apart,  since  by  their  interaction  potassium  hypochlorite 
and  potassium  chloride  would  be  formed  (cf.  p.  189).  In  the 
Castner-Kellner  apparatus  (Fig.  62),  which  serves  for  making  either 
potassium  or  sodium  hydroxide,  the  two  end  compartments  are 
filled  with  potassium  chloride  solution  (or  brine)  and  contain  the 
graphite  anodes.  The  central  compartment  contains  potassium 


864  COLLEGE   CHEMISTRY 

hydroxide  solution  and  the  iron  cathode.  The  positive  current 
enters  by  the  anodes,  and  the  chlorine  is  therefore  attracted  to  and 
liberated  upon  the  graphite:  2C1'  +  2©->  C12.  After  rising 
through  the  liquid  it  is  collected  for  the  manufacture  of  liquefied 
chlorine  or  of  bleaching  powder.  The  ions  of  potassium  (or  of 
sodium)  are  discharged  upon  a  layer  of  mercury  which  covers  the 
whole  floor  of  the  box,  and  the  free  metal  dissolves  in  the  mercury, 
forming  an  amalgam  (p.  352).  The  layer  of  mercury  extends 
beneath  the  partitions,  and  a  slight  rocking  motion  given  to  the  cell 
by  the  cam  (C)  causes  the  amalgam  to  flow  below  the  partition  into 
the  central  compartment.  Here  the  sodium  leaves  the  mercury  in 
the  form  of  sodium  ions  and  is  attracted  by  the  cathode.  Upon 
this,  hydrogen  from  the  water  is  discharged,  and  the  residual 
hydroxide-ion,  together  with  the  metal-ion,  constitutes  potassium 
or  sodium  hydroxide : 

2K*  +  2H'  +  2OH'  +  2Q->  2K*  +  20IT  +  H2. 

A  slow  influx  of  salt  solution  to  the  end  compartments,  and  overflow 
of  the  alkaline  solution  in  the  central  cell,  are  maintained.  The 
overflowing  liquid  contains  20  per  cent  of  the  alkali.  Since  there  is 
no  undecomposed  chloride  present  in  the  part  of  the  solution  which 
contains  the  hydroxide,  simple  evaporation  to  dryness  furnishes  the 
solid  alkali. 

Potassium  hydroxide  is  exceedingly  soluble  in  water,  and  conse- 
quently, instead  of  being  crystallized  from  solution,  the  molten 
residue  from  evaporation  is  cast  in  sticks.  The  hydroxide  is  highly 
deliquescent.  It  also  absorbs  carbon  dioxide  from  the  air,  giving 
potassium  carbonate.  Solutions  of  the  hydroxide  have  an  exceed- 
ingly corrosive  action  upon  the  flesh,  decomposing  it  into  a  slimy 
mass  by  hydrolyzing  the  albuminous  and  other  substances.  In 
solution,  the  base  is  highly  ionized,  furnishing  a  high  concentration 
of  hydroxidion.  Commercially  it  is  chiefly  employed  in  the  making 
of  soft  soap. 

Potassium  oxide  K2O  may  be  made  by  heating  potassium 
nitrate  with  potassium  in  a  vessel  from  which  air  is  excluded: 
KN03  +  5K  -»  3K2O  +  N.  It  interacts  violently  with  water,  giv- 
ing the  hydroxide.  When  exposed  to  the  air  it  unites  spontaneously 
with  oxygen,  and  a  yellow  peroxide  K204  is  formed. 


POTASSIUM  365 

Potassium  Chlorate  KClOy  —  The  preparation  of  this  salt 
by  interaction  of  potassium  chloride  with  calcium  chlorate,  has 
already  been  described  (p.  195).  It  is  also  made  by  electrolysis  of 
potassium  chloride  solution,  the  potassium  hydroxide  and  chlorine 
which  are  liberated  being  precisely  the  materials  required.  All  that 
is  necessary  is  to  use  a  warm,  concentrated  solution  and  to  provide 
for  the  mixing  of  the  materials  generated  at  the  electrodes.  The 
salt  crystallizes  out  when  the  solution  cools. 

Potassium  chlorate  crystallizes  in  monoclinic  plates.  It  melts  at 
about  351°,  and  at  a  temperature  slightly  above  this  the  visible 
liberation  of  oxygen  begins  (cf.  pp.  55,  196).  On  account  of  the 
ease  with  which  its  oxygen  is  liberated,  the  salt  is  employed  in  mak- 
ing fireworks  and  as  a  component,  along  with  antimony  trisulphide, 
of  the  heads  of  Swedish  matches.  It  is  also  used  in  medicine. 

Potassium  perchlorate  KC1O4,  formed  by  the  heating  of  the  chlorate 
(p.  196),  gives  white  crystals  belonging  to  the  rhombic  system. 

The  mode  of  preparing  potassium  bromate  KBrO3  and  potassium 
iodate  KIO3  has  already  been  described  (pp.  197,  198).  Potassium 
iodate  may  be  made  also  very  conveniently  by  melting  together 
potassium  chlorate  and  potassium  iodide  at  a  low  temperature. 
The  iodate  is  much  less  soluble  (see  Table)  than  the  chloride,  and 
the  mixture  may  be  separated  by  crystallization  from  water. 

Potassium  Nitrate  KNOy  —  The  formation  of  this  salt  in 
nature  and  its  mode  of  extraction  and  purification  have  already 
been  described  (p.  292).  This  source  of  supply  proved  insufficient* 
for  the  first  time,  during  the  Crimean  war  (1852-55),  and  a  method 
of  manufacture  from  Chili  saltpeter  (sodium  nitrate),  which  is  a 
much  cheaper  substance,  was  introduced.  Sodium  nitrate  and 
potassium  chloride  are  heated  with  very  little  water,  and  the 
sodium  chloride  produced  by  the  action,  which  is  a  reversible  one, 
is  by  far  the  least  soluble  of  the  four  salts  (see  Diagram,  p.  104).  On 
the  other  hand,  at  this  temperature,  the  potassium  nitrate  is  by  far 
the  most  soluble.  Hence  the  hot  liquid  drained  from  the  crystals 
contains  the  required  salt,  and  most  of  the  sodium  chloride  is  in  the 
form  of  a  precipitate.  If  the  solubility  curve  of  potassium  nitrate 
(p.  104)  is  examined,  it  will  be  seen  that  this  salt  is  but  slightly 
soluble  in  cold  water,  and  hence  most  of  it  is  deposited  when  the 
solution  cools.  The  crystals  are  mixed  with  little  sodium  chloride, 


366 


COLLEGE   CHEMISTRY 


for,  as  the  curve  shows,  common  salt  is  little  less  soluble  at  10°  than 
it  is  at  100°. 

Potassium  nitrate  gives  long  prisms  belonging  to  the  rhombic  sys- 
tem (Fig.  63).  It  melts  at  about  340°,  and  when  more  strongly 
heated  gives  off  oxygen,  leaving  potassium  nitrite 
(p.  299).  Although  it  does  not  form  a  hydrate, 
the  crystals  inclose  small  portions  of  the  mother- 
liquor,  and  consequently  contain  both  water  and 
impurities.  When  heated,  the  crystals  fly  to 
pieces  explosively  (decrepitate),  on  account  of  the 
vaporization  of  this  water.  Many  substances  which 
form  large  crystals  and  do  not  melt  at  a  low  tem- 
perature, behave  in  the  same  way  and  for  the  same 
reason.  In  consequence  of  this,  the  purest  salt 
is  made  by  violent  stirring  of  the  solution  during 
the  operation  of  crystallization,  the  result  being 
the  formation  of  a  crystal-meal. 

FIG.  63.  Potassium  nitrate  is  used  chiefly  in  the  manu- 

facture of  gunpowder,  which  contains  75  per  cent 
of  the  highly  purified  salt.  The  other  components  are  10  per  cent 
of  sulphur,  14  per  cent  of  charcoal,  and  about  1  per  cent  of  water. 
The  ingredients  are  intimately  mixed  in  the  form  of  paste,  and  the 
material  when  dry  is  broken  up  and  sifted,  grains  of  different  sizes 
being  used  for  different  purposes.  The  chemical  action  which 
takes  place  when  gunpowder  is  fired  in  an  open  space  gives  chiefly 
potassium  sulphide,  carbon  dioxide,  and  nitrogen: 


2KNO   +  3C  +  S 


K2S  +  3C02  +  Na. 


The  explosion  occurring  in  firearms  follows  a  much  more  complex 
course,  and  half  of  the  solid  product  is  said  to  be  potassium  carbonate. 
The  pressure,  at  the  temperature  of  the  explosion,  if  the  gases  could 
be  confined  within  the  volume  originally  occupied  by  the  gunpowder, 
would  reach  about  forty-four  tons  per  square  inch.  In  recent  years 
common  gunpowder  has  been  displaced  largely  by  smokeless  powder, 
of  which  substances  related  to  gun-cotton  (pp.  294,  301)  are  the 
chief  components. 

Potassium  Carbonate  K2CO3» — This  salt  is  manufactured  from 
potassium  chloride,  from  the  Stassfurt  deposits.     The  chloride  is 


POTASSIUM  867 

heated  with  magnesium  carbonate  (magnesite),  water,  and  carbon 
dioxide  under  pressure: 

2KC1  +  3MgC03  +  C02  +  5H20  ->  2KHMg(C03)2,4H20  +  MgCl2. 

The  hydrated  mixed  salt  is  separated  from  the  liquid  containing 
magnesium  chloride  and  decomposed  by  heating  with  water  at  120°. 
The  product  is  a  solution  of  potassium  carbonate,  from  which  the 
precipitated  magnesium  carbonate  is  removed  by  filtration.  In 
some  districts  potassium  carbonate  is  still  extracted  from  wood- 
ashes. 

This  salt  is  usually  sold  in  the  form  of  an  anhydrous  powder  (m.-p. 
over  1000°).  When  crystallized  from  water  it  gives  a  hydrate 
2K2CO3,3H2O.  It  is  deliquescent.  Its  aqueous  solution,  like  that 
of  sodium  carbonate  (cf.  p.  353),  has  a  marked  alkaline  reaction. 
The  commercial  name  of  the  substance  is  pearl  ash.  It  is  used  in 
making  soft  soap  and  hard  glass.  It  is  also  employed,  by  interaction 
with  acids,  in  making  salts  of  potassium. 

Potassium  Cyanide  KNC.  —  This  salt  is  made  by  heating 
together  dry  potassium  ferrocyanide  (q.v.)  and  potassium  carbonate. 
The  ferrocyanide  acts  as  if  it  were  a  mixture  of  potassium  cyanide 
and  ferrous  cyanide:  K4Fe(CN)6  -»  4KCN  +  Fe(CN)2.  The  latter, 
by  interaction  with  the  potassium  carbonate,  would  give  potassium 
cyanide  and  ferrous  carbonate,  but  this  in  turn,  is  decomposed  by 
heat  into  ferrous  oxide,  which  is  insoluble,  and  carbon  dioxide: 

K4Fe(CN)6  +  K2CO3  -*  6KCN  +  FeO  +  CO2. 

When  the  residue  is  extracted  with  water,  only  the  potassium 
cyanide  dissolves,  and  it  is  easily  crystallized  in  pure  form  from  the 
solution. 

Potassium  cyanide  is  extremely  soluble  in  water,  and  is  therefore 
deliquescent.  Its  poisonous  qualities  are  equal  to  those  of  hydro- 
cyanic acid.  The  acid  is  so  feeble  as  to  be  liberated  both  by  the 
moisture  and  by  the  carbon  dioxide  of  the  air,  and  hence  this  salt 
always  has  an  odor  of  hydrocyanic  acid.  Potassium  cyanide  is  used 
in  electroplating  (q.v.),  and  in  extracting  gold  (q.v.)  from  its  ores. 

The  preparation  of  potassium  cyanate  KCNO,  a  white,  easily 
soluble  salt,  and  of  potassium  thiocyanate  KCNS,  a  white,  deliquescent 
salt,  have  already  been  described  (p.  336). 


368  COLLEGE   CHEMISTRY 

The  Sulphate  and  Bisulphate.  —  Potassium  sulphate  K2S04 
is  a  constituent  of  several  double  salts  found  in  the  Stassfurt  de- 
posits. It  is  extracted  from  schoenite  MgSO4,K2S04,6H2O  and 
kainite  MgS04,MgCl2,K2SO4,6H20.  The  former  is  treated  with 
potassium  chloride  and  comparatively  little  water,  whereupon  the 
relatively  insoluble  potassium  sulphate  crystallizes  out,  and  the 
magnesium  chloride  remains  in  the  mother-liquor.  The  crystals 
belong  to  the  rhombic  system,  contain  no  water  of  crystallization, 
and  melt  at  1066°.  This  salt  is  employed  in  preparing  alum  (q.v.) 
and  is  much  used  as  a  fertilizer.  Since  plants  take  up  so- 
lutions through  their  cell  walls,  they  can  absorb  soluble  compounds 
only.  They  are,  therefore,  dependent,  for  the  potassium  com- 
pounds which  they  require,  upon  the  weathering  out  of  soluble 
potassium  compounds  from  insoluble  silicates  containing  potas- 
sium (p.  348)  found  in  the  soil.  The  weathering  takes  place  too 
slowly  to  furnish  a  sufficient  supply  for  many  crops,  particularly 
that  of  the  sugar-beet.  Hence  potassium  sulphate  is  mixed 
directly  with  the  soil. 

Potassium-hydrogen  sulphate  (bisulphate)  KHS04  is  made  by  the 
action  of  sulphuric  acid  upon  potassium  sulphate:  K2SO4  +  H2S04  — » 
2KHS04.  It  crystallizes  from  water,  in  which  it  is  very  soluble,  in 
tabular  crystals.  Its  properties  are  similar  to  those  of  sodium 
bisulphate  which  have  already  been  described  (p.  265). 

Sulphides  of  Potassium.  —  By  the  treatment  of  a  solution  of 
potassium  hydroxide  with  excess  of  hydrogen  sulphide,  a  solution  of 
potassium-hydrogen  sulphide  is  obtained.  Evaporation  of  the  solu- 
tion gives  a  deliquescent,  solid  hydrate  2KHS,H2O.  When  the 
solution,  before  evaporation,  is  treated  with  an  equivalent  amount 
of  potassium  hydroxide,  and  the  water  is  driven  off,  the  sulphide 
K2S  remains  behind  (cf.  p.  253) : 

KHS  +  KOH  <=±  K2S  +  H20. 

Considerable  amounts  of  sulphur  can  be  dissolved  in  solutions  of 
either  of  these  sulphides.  By  evaporation  of  the  resulting  yellow 
liquids,  various  polysulphides  have  been  obtained.  To  some  of  these 
have  been  ascribed  the  formulae  K,S3,  K2S4,  and  K2S5  (cf.  p.  255). 
Similar  substances  are  produced,  as  a  result  of  the  liberation  and 


POTASSIUM  :!t;<i 

recombination  of  sulphur,  when  the  solutions  are  exposed  to  the 
oxidizing  action  of  the  air : 

2KHS  +  O2  ->  2KOH  +  2S. 

Properties  of  Potassium- ion  K':  Analytical  Reactions, — The 

positive  ionic  material  of  the  potassium  salts  is  a  colorless  substance. 
It  unites  with  all  negative  ions,  and  most  of  the  resulting  compounds 
are  fairly  soluble.  For  its  recognition  we  add  solutions  containing 
those  ions  which  give  with  it  the  least  soluble  salts.  Thus,  with 
chloroplatinic  acid  H2PtCl6  it  gives  a  yellow  precipitate  of  potassium 
chloroplatinate  K2PtCltt.  Since  nearly  one  part  of  this  salt  dissolves 
in  100  parts  of  water,  the  test  is  far  from  being  a  delicate  one.  Picric 
acid  (p.  294)  gives  potassium  picrate  KC6H2(NO2)3O,  which  is  much 
less  soluble  in  water  (0.4  parts  in  100  at  15°).  Perchloric  acid  and 
hydrofluosilicic  acid  likewise  give  somewhat  insoluble  salts  of  potas- 
sium. Potassium-hydrogen  tartrate  KHC4H4O0  is  precipitated  by 
the  addition  of  tartaric  acid  to  a  sufficiently  concentrated  solution 
of  a  potassium  salt.  The  neutral  tartrate  K2C4H406  is  much  more 
soluble.  The  latter  may  be  obtained  by  treating  the  precipitate 
with  a  solution  of  potassium  hydroxide.  Addition  of  an  acid  to 
this  solution  causes  reprecipitation  of  the  bitartrate. 

A  much  more  delicate  test  for  the  recognition  of  a  potassium  com- 
pound consists  in  the  examination  by  means  of  the  spectroscope  of 
the  light  given  out  by  a  Bunsen  flame,  in  which  a  little  of  the  salt  is 
held  upon  a  platinum  wire.  When  the  amount  of  potassium  is  con- 
siderable, and  no  other  substance  which  would  likewise  color  the 
flame  is  present  to  mask  the  effect,  the  violet  tint  is  recognizable 
by  the  eye. 

Rubidium  and  Caesium.  —  In  1860  Bunsen  discovered  several 
new  lines  in  the  spectrum  given  by  materials  derived  from  the  salts 
in  Durkheim  mineral  water.  Two  new  elements  of  the  alkali  group 
were  found  to  cause  their  presence,  and  were  named,  from  the  colors 
of  the  lines  which  they  gave,  rubidium  (red)  and  caesium  (blue). 
Rubidium  is  obtainable  with  relative  ease  from  the  mother-liquors 
of  the  Stassfurt  works. 

The  metals  may  be  obtained  by  heating  their  hydroxides  with 
magnesium  powder.  The  hydroxides  of  these  two  elements  are 
more  active  than  potassium  hydroxide.  Their  salts  are  very  much 
like  those  of  potassium. 


370  COLLEGE   CHEMISTRY 

AMMONIUM. 

The  compounds  of  ammonium  claim  a  place  with  those  of  the 
alkali  metals  because  in  aqueous  solution  they  give  ammonium-ion 
NH4",  a  substance  which  in  its  behavior  closely  resembles  potassium- 
ion.  Some  of  the  special  properties  peculiar  to  ammonium  com- 
pounds, and  particularly  the  properties  of  ammonium  hydroxide 
NH4OH,  have  been  discussed  in  detail  already  (pp.  282-283). 

Salts  of  Ammonium.  —  Ammonium  chloride  NH4C1,  known  com- 
mercially as  salammoniac,  like  all  the  other  compounds  of  ammonium, 
is  prepared  from  the  ammonia  dissolved  by  the  water  used  to  wash 
illuminating-gas  (p.  281).  It  is  purified  by  sublimation,  and  then 
forms  a  compact  fibrous  mass.  When  heated  to  350°  it  volatilizes 
and  is  almost  completely  dissociated  into  ammonia  and  hydrogen 
chloride  at  this  temperature  (p.  283). 

Ammonium  nitrate  NH4NO3  is  a  white  crystalline  salt  which  may 
be  made  by  the  interaction  of  ammonium  hydroxide  and  nitric  acid. 
When  heated  gently  it  decomposes,  giving  nitrous  oxide  and  water 
(p.  300).  It  is  used  as  an  ingredient  in  fireworks  and  explosives. 

When  ammonium  hydroxide  is  treated  with  excess  of  carbon 
dioxide  the  solution  gives,  on  evaporation,  ammonium  bicarbonate 
NH4HCO3.  This  is  a  white  crystalline  salt  which  is  fairty  stable  at 
the  ordinary  temperature.  It  has,  however,  a  faint  odor  of  am- 
monia, and  its  dissociation  becomes  very  rapid  when  slight  heat  is 
applied.  When  a  solution  of  this  salt  is  treated  with  ammonium 
hydroxide,  the  neutral  carbonate  (NH4)2C03  is  formed.  But  this  salt, 
when  left  in  an  open  vessel,  loses  ammonia  very  rapidly,  and  leaves 
the  bicarbonate  behind. 

Ammonium  thiocyanate  NH4NCS  (cf.  p.  179)  is  a  white  salt  which 
finds  some  application  in  analysis. 

Ammonium  sulphate  (NH4)2SO4  is  a  white  salt  which  is  used  chiefly 
as  a  fertilizer.  By  electrolysis  of  a  concentrated  solution  of  the 
bisulphate  NH4HS04,  ammonium  persulphate  (NH4)2S2OS,  which  is 
less  soluble,  is  formed  and  crystallizes  out  (cf.  p.  267). 

Solutions  of  ammonium-hydrogen  sulphide  NH4HS  and  ammonium 
sulphide,  (NH4)2S,  made  by  passing  hydrogen  sulphide  gas  into 
ammonium  hydroxide,  are  much  used  in  analysis.  The  sulphide  is 
almost  completely  hydrolyzed  by  water  into  the  acid  sulphide  and 


AMiMUNir.M 

ammonium  hydroxide,  its  behavior  being  like  that  of  sodium 
sulphide  (p.  253) : 

2  NH3  +  H2S  +±  (NH4)2S  fc?  2  NH/  +  S"  L      ™, 
H20±=;      pB'  +  H-r*1 

It  is  used  for  the  precipitation  of  sulphides,  such  as  zinc  sulphide, 
which  are  insoluble  in  water.  Although  the  S"  ions  are  not  numer- 
ous at  any  moment,  disturbance  of  the  equilibrium  by  their  removal, 
when  they  pass  into  combination,  causes  displacements  which  result 
in  the  generation  of  a  continuous  supply.  The  liquid  smells  strongly 
of  ammonia  and  hydrogen  sulphide,  on  account  of  the  dissociation 
of  the  parent  molecules  by  reversal  of  the  above  equilibria. 

The  solutions,  when  pure,  are  colorless.  They  dissolve  free  sulphur, 
giving  yellow  polysulphides  similar  to  those  of  potassium  (p.  368). 
The  same  yellow  substances  are  also  obtained  by  gradual  oxidation 
of  ammonium  sulphide  when  the  solution  of  this  salt  is  allowed  to 
stand  in  a  bottle  from  which  the  air  is  imperfectly  excluded. 

Ammonium  Amalgam.  —  When  a  salt  of  ammonium  is  decom- 
posed by  electrolysis  the  NH4",  upon  its  discharge,  ordinarily  gives 
ammonia  and  hydrogen,  and  no  substance  NH4  is  obtained.  If, 
however,  a  pool  of  mercury  is.  used  as  the  negative  electrode,  the  NH4 
forms  an  amalgam  with  it,  and  there  seems  to  be  no  doubt  that  this 
substance  is  actually  present  in  solution  in  the  mercury.  While  the 
amalgam  is  being  formed  it  swells  up  and  gives  off  the  decomposi- 
tion products  above  mentioned,  so  that  the  existence  of  the  sub- 
stance is  only  temporary.  The  same  material  may  be  obtained  by 
putting  sodium  amalgam  into  a  strong  solution  of  a  salt  of  ammonium. 
The  action  is  a  displacement  of  one  ion  by  another  (p.  243) : 

Na(dslvd  in  mercury)  +  NH4*  ->  NH4(dslvd  in   mercury)  +  Na*. 

This  behavior  is  interesting  since  it  is  in  harmony  with  the  idea  that 
ammonium,  if  it  could  be  isolated,  would  have  the  properties  of  a 
metal.  Substances  other  than  metals  are  not  miscible  with 
mercury. 

Ammonium-ion  jVJBT/r  Analytical  tteactions.  —  Ionic  ammo- 
nium is  a  colorless  substance.  It  unites  with  negative  ions,  giving 
salts,  which,  in  the  majority  of  cases,  are  soluble.  Ammonium 
chloroplatinate  (NH4)2PtCl6,  and  to  a  less  extent  ammonium-hydro- 


372  COLLEGE  CHEMISTRY 

gen  tartrate  NH4HC4H4Oe,  are  insoluble  compounds,  and  their  pre- 
cipitation is  used  as  a  test.  The  surest  means  of  recognizing  ammo- 
nium compounds,  however,  consists  in  adding  a  soluble  base  to  the 
substance  (cf.  p.  283).  The  ammonium  hydroxide,  which  is  thus 
formed,  gives  off  ammonia,  and  the  latter  may  be  detected  by  its 
odor. 

Exercises. —  1.   What  kind  of  metals  will,  in  general,  interact 
with  solutions  of  bases  (cf.  p.  356)? 

2.  Why  should  a  mixture  of  potassium  chlorate  and  antimony 
trisulphide  be  explosive? 

3.  How  should  you  set  about  making,  (a)  a  borate  of  potassium, 
(b)  potassium  pyrophosphate,  (c)  ammonium  nitrite,  (d)  ammonium 
chlorate,  (e)  ammonium  iodide? 


CHAPTER  XXXIV 

SODIUM   AND   LITHIUM.     IONIC   EQUILIBRIUM   CONSIDERED 
QUANTITATIVELY 

SODIUM  chloride  forms  more  than  two-thirds  of  the  solid  matter 
dissolved  in  sea-water,  and  the  great  salt  deposits  are  largely  com- 
posed of  it.  Sea-plants  contain  sodium  salts  of  organic  acids,  just 
as  land-plants  contain  potassium  salts.  Chili  saltpeter,  cryolite, 
and  albite  (a  soda  feldspar)  are  important  minerals. 

Preparation.  —  Sodium  was  first  made  by  Davy  (1807)  by 
electrolysis  of  moist  sodium  hydroxide.  It  is  manufactured  by 
Castner's  process,  which  is  used  also  for  potassium  (p.  361),  and 
by  the  electrolysis  of  fused  sodium  hydroxide  by  a  method  likewise 
invented  by  Castner.  In  the  latter  case  the  negative  electrode  pro- 
jects through  the  bottom  of  the  iron  vessel  containing  the  fused 
hydroxide.  This  electrode  is  surrounded  by  a  wire-gauze  partition, 
which  is  surmounted  by  a  bell-shaped  vessel  of  iron.  The  positive 
electrode  is  an  iron  cylinder  surrounding  the  gauze.  The  sodium 
and  hydrogen  liberated  at  the  cathode,  being  lighter  than  the  fused 
mass,  ascend  into  the  iron  vessel,  under  the  edge  of  which  the  hydro- 
gen escapes.  Oxygen  is  set  free  at  the  anode. 

Properties.  —  Sodium  is  a  soft,  shining  metal,  melting  at  95.6° 
and  boiling  at  742°.  The  vapor  is  a  monatomic  gas.  The  general 
chemical  properties  have  already  been  given  (p.  360).  The  metal 
unites  with  hydrogen  to  form  a  hydride  NaH,  which  resembles  potas- 
sium hydride  (p.  361).  The  amalgam  with  mercury  when  it  contains 
more  than  a  small  amount  of  sodium,  is  solid,  and  probably  contains 
one  or  more  compounds  of  the  two  elements.  This  amalgam  is 
often  used  instead  of  the  metal  sodium,  since  the  dilution  or  combina- 
tion with  mercury  makes  the  interactions  of  the  metal  more  easily 
controllable.  Sodium  is  used  in  the  manufacture  of  many  complex 
carbon  compounds  which  are  employed  as  drugs  and  dyes. 

373 


374  COLLEGE   CHEMISTRY 

Sodium  Chloride  Nad.  —  Common  salt  is  obtained  from  the 
salt  deposits  of  Stassfurt,  Reichenhall  (near  Salzburg),  in  Cheshire, 
at  Syracuse  and  Warsaw  in  New  York,  at  Salina  in  Kansas,  in  Utah, 
California,  and  many  other  districts.  Natural  brines  are  obtained 
from  wells  in  various  parts  of  the  world.  Since  the  salt  can  seldom 
be  used  directly,  on  account  of  impurities  which  it  contains,  it  is 
purified  by  recrystallization  from  water.  Natural  brines,  which  are 
sometimes  dilute,  are  often  concentrated  by  dripping  over  extensive 
ricks  composed  of  twigs.  When  the  resulting  brine  is  allowed  to 
evaporate  slowly  by  the  help  of  the  sun's  heat,  large  crystals,  sold  as 
"  solar  salt,"  are  obtained.  By  the  use  of  artificial  heat  and  stirring, 
smaller  crystals  of  greater  purity  can  be  secured.  Salt  intended  for 
table  use  must  be  freed  from  the  traces  of  magnesium  chloride  (q.v.) 
present  in  the  original  brine  or  deposit,  for  this  impurity  causes  it  to 
absorb  moisture  more  vigorously  from  the  air.  The  purest  salt  for 
chemical  purposes  is  precipitated  from  a  saturated  solution  of  salt 
by  leading  into  it  hydrogen  chloride  gas.  Explanation  of  this  effect 
will  be  given  presently  (see  pp.  385-387). 

Common  salt  crystallizes  in  cubes,  the  faces  of  which  are  usually 
hollow.  The  crystals  decrepitate  (p.  366)  when  heated,  and  melt 
at  about  820°.  Common  salt  is  the  source  of  all  sodium  compounds, 
with  the  exception  of  the  nitrate.  From  it  come  also  most  of  the 
chlorine  and  hydrogen  chloride  used  in  commerce. 

The  Hydroocide  and  Oxides.  —  Sodium  hydroxide  NaOH,  called 
also,  colloquially,  caustic  soda,  is  prepared  by  the  action  of  slaked 
lime  upon  sodium  carbonate,  and  by  the  electrolysis  of  a  solution  of 
sodium  chloride,  in  both  cases  precisely  as  is  potassium  hydroxide 
(p.  363).  Sodium  hydroxide  is  a  highly  deliquescent  substance. 
Its  general  chemical  properties  are  identical  with  those  of  potassium 
hydroxide.  It  is  used  in  the  manufacture  of  soap,  in  the  preparation 
of  paper  pulp,  and  in  many  other  chemical  industries. 

Sodium  peroxide  Na202  is  made  by  heating  sodium  at  300-400°  in 
air  which  has  been  freed  from  carbon  dioxide.  This  oxide  is  the 
sodium  salt  of  hydrogen  peroxide.  When  thrown  into  water  it 
decomposes  in  part,  in  consequence  of  the  heat  developed,  giving 
sodium  hydroxide  and  oxygen.  With  careful  cooling,  however^ 
much  .of  it  can  be  dissolved.  By  interaction  with  acids  it  yields 
hydrogen  peroxide  (p.  211).  Sodium  peroxide  is  now  used  com- 


^ 

SODIUM 


375 


mercially  for  oxidizing  and  bleaching.     The  ordinary  sodium  oxide 
Na2O  is  made  in  the  same  way  as  is  potassium  oxide  (p.  364). 

The  Nitrate  and  Nitrite.  —  The  occurrence  and  purification 
of  sodium  nitrate  NaNO3  have  already  been  described  (p.  292).  Its 
crystals  are  of  rhombohedral  form  (Fig.  9,  p.  9).  This  salt  is  one 
of  the  best  of  fertilizers,  since  it  furnishes  to  plants  the  nitrogen 
which  they  require  in  a  very  easily  absorbed  form.  It  is  used  also 
in  the  manufacture  of  potassium  nitrate,  and  of  nitric  acid. 

Sodium  nitrite  NaN02  is  formed  by  heating  sodium  nitrate  with 
metallic  lead  and  recrystallizing  the  product  (p.  299). 

Manufacture  of  Sodium  Carbonate.  —  Natural  sodium  carbon- 
ate is  found  in  Egypt  and  in  other  parts  of  the  world.  At  Owen's 


FIG.  64. 

Lake,  California,  it  is  secured  by  solar  evaporation  of  the  water. 
The  sesquicarbonate  Na2CO3,NaHC03,2H2O,  being  the  least  sol- 
uble of  the  carbonates  of  sodium,  is  the  one  deposited.  Locally, 
small  quantities  of  sodium  carbonate  are  still  made  by  the  burning 
of  sea-weed.  The  substance  is  manufactured  from  sodium  chloride 
in  two  ways,  namely  by  the  Le  Blanc  process  and  by  the  Solvay 
process. 

The  Le  Blanc  process  (1791)  involves  three  chemical  actions.  In 
the  first  place,  sodium  chloride  is  treated  with  an  equivalent  amount 
of  sulphuric  acid  in  a  large  cast-iron  or  earthenware  pan.  The 


376  COLLEGE    CHEMISTRY 

bisulphate  thus  produced  (cf.  p.  117),  together  with  the  unchanged 
sodium  chloride,  is  raked  out  on  to  the  hearth  of  a  reverberatory  * 
furnace  (Fig.  64)  and  heated  more  strongly,  while  being  continually 
worked  by  means  of  rakes,  until  the  action  is  completed: 

NaCl  +  NaHSO4  <=>  Na,SO4  +  HC1 1  • 

The  product  of  this  treatment  is  called  salt-cake.  The  hydrogen 
chloride,  which  is  liberated  in  both  stages,  passes  through  towers 
containing  running  water  in  which  it  is  absorbed.  The  second  and 
third  actions  which  follow  are  conducted  in  one  operation.  They 
consist  in  the  reduction  of  the  sodium  sulphate  by  means  of  powdered 
coal  and  the  interaction  of  the  resulting  sulphide  of  sodium  with 
chalk  or  powdered  limestone : 

Na,S04  +  2C  ->  Na2S  +  2CO2, 
Na,,S  +  CaCO3  ->  Na2C03  +  CaS. 

In  the  less  modern  factories  the  salt-cake,  limestone,  and  coal  are 
stirred  upon  the  hearth  of  a  reverberatory  furnace  and  worked  by 
hand.  The  material  is  finally  collected  into  balls,  and  the  end  of  the 
action  is  recognized  by  the  fact  that  bubbles  of  carbon  monoxide 
begin  to  force  their  way  to  the  surface  and  cause  little  jets  of  blue 
flame.  The  gas  is  produced  by  the  action  of  the  coal  upon  the 
calcium  carbonate,  excess  of  both  of  these  substances  being  present: 
CaCO3  +  C  ->  CaO  +  2CO.  The  production  of  this  gas  gives  a 
porous  texture  to  the  material,  which  facilitates  the  solution  of  the 
sodium  carbonate  in  the  final  stage.  The  porous  product  is  called 
black-ash.  In  modern  factories  hand  labor  is  saved  by  giving  the 
black-ash  furnace  the  form  of  a  rotating  cylinder,  in  which  pro- 
jections from  the  walls  assist  in  bringing  about  complete  mixing  of 
the  materials  during  the  action. 

The  black-ash  varies  very  much  in  composition.  It  commonly 
contains  45  per  cent  of  sodium  carbonate,  30  per  cent  of  calcium 
sulphide,  10  per  cent  of  calcium  oxide,  and  a  number  of  other  prod- 
ucts and  impurities. 

Calcium  sulphide  is  not  very  soluble  in  water,  and  is  but  slowly 
hydrolyzed  by  it  (p.  254),  especially  when  calcium  hydroxide  is 
present.  The  sodium  carbonate  is  therefore  extracted  from  the 

*  So  called  because  the  heated  gases  from  the  fire  are  deflected  by  the  roof 
and  play  upon  the  materials  spread  on  the  bed  of  the  furnace, 


SODIUM  377 

black-ash  by  a  systematic  treatment  of  the  ash  with  water.  The 
ash  is  placed  in  a  series  of  vessels  at  different  levels,  and  a  stream  of 
water  (30-40°)  flows  from  one  vessel  to  another,  until,  when  it  issues 
from  the  last,  it  is  completely  saturated  with  sodium  carbonate. 
When  the  material  in  the  first  of  the  vessels  has  been  exhausted,  the 
water  is  allowed  to  enter  the  second  vessel  directly,  and  a  vessel 
containing  fresh  black-ash  is  added  at  the  lower  end  of  the  series. 
In  this  way  the  most  nearly  exhausted  ash  comes  in  contact  with 
pure  water,  which  is  in  the  best  position  to  dissolve  the  remaining 
sodium  carbonate  rapidly,  while  the  fresh  black-ash  encounters  a 
solution  already  almost  at  the  point  of  saturation. 

The  saturated  solution  is  evaporated  in  shallow  pans  placed  on  the 
flues  of  the  furnaces,  and  the  monohydrate  Na/X^H-jO,  which 
crystallizes  from  the  hot  liquid,  is  raked  out  and  dried  by  heat, 
leaving  calcined  soda.  When  this  material  is  recrystallized  from 
water  and  is  allowed  to  deposit  itself  from  the  solution  at  the  ordi- 
nary temperature,  the  decahydrate  Na2CO3,10H2O,  soda  crystals  or 
washing  soda,  appears. 

The  solid  residue  from  the  extraction  of  the  black-ash  is  known  as 
tank  waste,  and  contains  35-55  per  cent  of  calcium  sulphide.  This 
material  contains  the  sulphur  of  the  original  sulphuric  acid,  and  its 
treatment  involves  a  problem  of  some  difficulty.  If  it  is  dumped 
near  the  factory  the  sulphur  is  lost,  and  by  slow  weathering  yellow 
solutions  containing  polysulphides  flow  from  the  decomposing  heap 
into  the  streams,  and  offensive  odors  of  hydrogen  sulphide  fill  the 
air.  The  most  effective  process  for  the  recovery  of  the  sulphur  and 
consequent  abatement  of  this  nuisance  is  that  of  Chance.  The 
product  is  arranged  in  a  series  of  cylinders  through  which  is  passed 
carbon  dioxide  from  a  kiln  of  special  form.  The  hydrogen  sulphide 
liberated  in  the  first  cylinder  forms  the  acid  sulphide  with  the 
material  contained  in  the  second: 

2CaS  +  CO2  +  H2O  -»  Ca(SH)2  +  CaC03. 

The  further  action  of  the  carbon  dioxide  on  this  product  gives,  finally, 
a  mixture  of  gases  containing  a  larger  proportion  of  hydrogen  sul- 
phide. By  burning  this  mixture  with  a  limited  supply  of  air,  the 
sulphur  is  then  secured  in  free  condition: 

2H2S  +  02-»2H2O  +2S. 
About  70,000  tons  of  sulphur  are  thus  recovered  annually. 


378  COLLEGE   CHEMISTRY 

The  Solvay,  or  ammonia-soda  process  (1860),  is  a  serious  rival  of  the 
Le  Blanc  process.  It  differs  from  the  latter  by  involving  almost 
nothing  but  ionic  actions.  A  solution  of  salt  containing  ammonia 
and  warmed  to  40°  fills  a  tower  divided  by  a  number  of  perforated 
partitions.  Carbon  dioxide,  which  is  forced  in  below,  makes  its  way 
up  through  the  liquid.  The  ammonium  bicarbonate  formed  by  its 
action  undergoes  double  decomposition  with  the  salt,  and  sodium 
bicarbonate  which  is  precipitated  settles  upon  the  partitions: 

NaCl  +  NH4HC03  ±;NaHCO,  J  +  NH4C1,  or  HC03'  +  Na'«=±NaHCO, J. 

The  solid  sodium  bicarbonate,  after  being  freed  from  the  liquid,  is 
heated  strongly  and  leaves  behind  sodium  carbonate: 

2NaHCO3  <-  Na.CO,  +  H2O  |  +  CO2 1. 

The  carbon  dioxide  which  is  liberated  passes  through  the  operation 
once  more.  The  mother-liquor  from  the  sodium  bicarbonate  con- 
tains ammonium  chloride.  This  is  decomposed  by  heating  with 
quicklime,  and  the  ammonia  which  is  thus  obtained  is  available  for 
the  treatment  of  another  batch.  The  supply  of  carbon  dioxide  is 
generated  in  lime-kilns  of  special  form.  The  lime  produced  in  these 
kilns  serves  for  the  liberation  of  the  ammonia. 

This  process  is  cheaper  than  that  of  Le  Blanc,  and  furnishes  a 
much  purer  product.  The  latter  process  continues  to  be  used,  how- 
ever, although  not  in  the  United  States,  because  of  the  hydrochloric 
acid  which  is  produced  at  the  same  time.  This  finds  remunerative 
application  in  the  liberation  of  its  chlorine  for  the  manufacture  of 
bleaching  powder. 

Properties  of  the  Carbonate  and  Bicarbonate.  —  The  com- 
mon form  of  sodium  carbonate  consists  of  large  monoclinic  crystals 
of  the  decahydrate  Na2CO3,10H2O.  This  substance  has  a  fairly  high 
aqueous  tension,  and  loses  nine  of  the  ten  molecules  of  water  which 
it  contains  when  it  is  exposed  in  an  open  vessel  (p.  83).  When 
warmed  it  melts  at  35. 2°,  giving  a  solution  of  sodium  carbonate  in 
water.  The  residue  from  evaporation,  above  35.2°,  is  the  mono- 
hydrate  Na2CO3,H2O.  At  higher  temperatures,  or  in  a  dry  atmos- 
phere (p.  83),  this  in  turn  can  be  completely  dehydrated.  In 
aqueous  solution,  sodium  carbonate  is  hydrolyzed,  and  shows  a 
marked  alkaline  reaction  (p.  353).  The  compound  is  used  in  large 
amounts  for  the  manufacture  of  glass  and  soap,  and  is  applied  in 


SODIUM  379 

innumerable  ways  in  the  scientific  industries  for  purposes  akin  to 
cleansing. 

Nearly  all  the  familiar  compounds  of  sodium  are  formed  in  the 
course  of  one  or  other  of  the  processes  by  which  sodium  carbonate  is 
manufactured,  or  are  made  by  the  treatment  of  sodium  carbonate  or 
sodium  hydroxide  with  acids. 

Sodium  bicarbonate  NaHC03  is  formed  in  the  Solvay  process 
(p.  378).  It  can  be  prepared  in  a  state  of  purity  by  passing  carbon 
dioxide  over  the  decahydrate  of  sodium  carbonate: 

NaaCO^lOHaO  +  CO2  <=±  2NaHCO3  +  9H20. 

This  action  is  reversible  (cf.  p.  174),  and  sodium  bicarbonate  shows, 
even  in  the  cold,  an  appreciable  tension  of  carbon  dioxide.  The 
aqueous  solution  is  neutral  to  phenolphthalein,  on  account  of  the 
small  degree  of  ionization  of  the  ion  HCO3'.  The  salt  is  used  in 
the  manufacture  of  baking  powder  and  in  medicine. 


Other  Salts  of  Sodium.  —  Anhydrous  sodium  sulphate 

(thenardite)  is  found  in  the  salt  layers.  The  same  salt  is  contained 
in  mineral  waters,  such  as  those  of  Friedrichshall  and  Karlsbad.  It 
is  formed  in  connection  with  the  manufacture  of  nitric  acid  from 
sodium  nitrate,  and  as  an  intermediate  product  in  the  making  of 
sodium  carbonate. 

The  decahydrate  of  sodium  sulphate  Na2SO4,10H2O  (Glauber's  salt) 
forms  large  monoclinic  crystals  which  give  up  ten  molecules  of  water 
when  kept  in  an  open  vessel.  When  heated  the  crystals  melt  at 
32.4°,  and  are  resolved  into  the  sulphate  and  water.  The  solubili- 
ties of  the  hydrate  and  anhydrous  substance  are  shown  in  Fig.  41 
(p.  106). 

Sodium  thiosulphate  Na2S2O3,5H20,  formerly  called  hyposulphite 
of  soda,  and  still  called  hypo  by  photographers,  is  made  by  boiling 
a  solution  of  sodium  sulphite  with  sulphur.  It  is  also  obtained  by 
boiling  sulphur  with  caustic  soda,  and  crystallizes  from  the  mixed 
solution  : 

4S  +  6NaOH  ->  Na2S2O3  +  2Na,>S  +  3H2O. 

When  heated  it  first  loses  the  water  of  hydration,  and  then  decom- 
poses, giving  sodium  sulphate,  which  is  the  most  stable  oxygen- 
sulphur  compound  of  sodium  (cf.  p.  266)  and  sodium  pentasulphide: 

4Na2S203  ->  3Na,S04  +  Na2S5. 


380  COLLEGE   CHEMISTRY 

From  the  latter,  four  unit-weights  of  sulphur  can  be  driven  by 
stronger  heating.  Sodium  thiosulphate  is  used  for  fixing  negatives 
in  photography  (q.v.),  and  by  bleachers  as  antichlor. 

Common  sodium  phosphate  is  a  dodecahydrate  of  the  secondary 
orthophosphate,  Na2HPO4,12H2O.  It  is  made  by  neutralization  of 
phosphoric  acid  with  sodium  carbonate.  Its  properties  have 
already  been  discussed  (pp.  311-312). 

Sodium  metaphosphate  NaPO3  is  used  for  bead  reactions  (p.  312). 

Sodium  tetraborate  Na^I^O^lOHgO  (borax)  forms  large,  trans- 
parent prisms.  When  heated  it  loses  water,  and  leaves  the  easily 
fusible  anhydrous  salt  in  glassy  form.  Its  sources  have  already 
been  discussed  under  borates  (p.  350).  It  is  used  as  an  ingredient 
in  glazes  for  porcelain,  in  soldering,  for  bead  reactions  (p.  350)  and 
for  preserving  food. 

Sodium  metasilicate  Na-jSiOg  (cf.  p.  346)  is  used  for  fire-proofing 
wood  and  other  materials,  and  for  preserving  eggs.  Sand  which 
is  moistened  with  it  and  pressed  in  molds,  forms,  after  baking,  a 
serviceable  artificial  stone. 

Properties  of  Sodium-ton  :  Analytical  Reactions.  —  Sodium- 
ion  Na*  is  a  colorless  ionic  material  which  unites  with  all  negative 
ions.  Practically  all  the  salts  so  formed  are  soluble  in  water.  The 
only  ones  which  can  be  precipitated  are  sodium  fluosilicate  Na^SiFg, 
made  by  the  addition  of  hydrofluosilicic  acid  to  a  strong  solution  of 
a  sodium  salt,  and  sodium-hydrogen  pyroantimoniate  Na^SbjO?, 
made  by  similar  addition  of  the  corresponding  potassium  salt.  All 
compounds  of  sodium  confer  a  yellow  color  on  the  Bunsen  flame, 
but  this  test  is  so  delicate  that  it  is  shown  by  the  traces  of  sodium 
contained  in  almost  all  substances. 

Lithium. — Lithium  occurs  in  lepidolite  (a  lithia  mica),  in  ambly- 
gonite,  and  in  other  rare  minerals.  Traces  of  compounds  of  the 
element  are  found  widely  diffused  in  the  soil,  and  are  taken  up  by 
plants,  particularly  tobacco  and  beets,  in  the  ashes  of  which  the 
element  may  be  detected  spectroscopically. 

The  metal  is  liberated  by  electrolysis  of  the  fused  chloride.  The 
specific  gravity  of  the  free  element  (0.53)  is  lower  than  that  of  any 
other  metal.  Lithium  not  only  floats  upon  water,  but  also  in  the 
petroleum  in  which  it  is  preserved. 


IONIC   EQUILIBRIUM,    CONSIDERED    QUANTITATIVELY         381 

The  metal  behaves  towards  water  and  oxygen  like  sodium  (p.  66). 
It  unites  directly  and  vigorously  with  hydrogen  (LiH),  nitrogen 
(Li3N),  and  oxygen  (Li2O),  forming  stable  compounds.  The  relative 
insolubility  (see  Table)  of  the  hydroxide  (LiOH),  the  carbonate 
(Li2CO3),  and  the  phosphate  (Li3PO4,2H20),  is  in  sharp  contrast  to 
the  easy  solubility  of  the  corresponding  compounds  of  the  other 
alkali-metals,  and  links  lithium  with  magnesium.  The  compounds 
of  lithium  give  a  bright-red  color  to  the  Bunsen  flame.  A  bright- 
red  and  a  somewhat  less  bright  orange  line  are  seen  in  the  spectrum. 
The  carbonate  is  used  in  medicine. 

IONIC  EQUILIBRIUM,  CONSIDERED  QUANTITATIVELY. 

In  view  of  the  predominance  of  ionic  actions  in  the  chemistry  of  the 
metals,  and  of  the  determinative  effect  of  ionic  equilibria  on  many 
actions,  it  is  essential  that  we  should  be  prepared  in  future  for  a 
more  exact  consideration  of  these  phenomena  than  we  have  hitherto 
attempted.  The  whole  basis  for  this  exact  consideration  has  already 
been  supplied,  and  only  more  specific  application  of  the  principles  is 
demanded.  The  basis  referred  to,  which  should  now  be  studied  as 
a  preliminary  to  what  follows,  is  contained  in,  (1)  the  discussion  of 
chemical  equilibrium  in  general  (pp.  174-185),  (2)  the  application  of 
the  same  principles  to  ionic  equilibrium  (p.  224),  and  (3)  the  illus- 
tration of  this  application  in  the  case  of  cupric  bromide  (pp.  233- 
236). 

Excess  of  One  Ion.  —  In  the  case  of  cupric  bromide,  we  showed 
that  increasing  the  concentration  of  the  bromide  ions  displaced  the 
equilibrium  by  favoring  the  union  of  the  ions  to  form  molecular 
cupric  bromide:  2Br'  +  Cu"  — » CuBr2.  This  we  speak  of  as  a 
repression  of  the  ionization  of  the  cupric  bromide.  Now,  if  the  sub- 
stance is  a  slightly  ionized  one,  like  a  weak  acid  or  a  weak  base,  the 
repression  of  the  ionization  through  the  formation  of  molecules  in 
this  way  may  remove  so  many  of  that  one  of  the  ions  which  is  not 
present  in  excess  (corresponding  to  the  Cu"  in  the  foregoing  illus- 
tration), that  the  mixture  will  no  longer  respond  to  tests  for  the  ion 
so  removed.  This  is  an  interesting  and  very  common  case.  The 
behavior  of  acetic  acid,  a  weak,  slightly  ionized  acid,  will  serve  as 
an  illustration. 


382  COLLEGE    CHEMISTRY 

In  normal  solution  (60  g.  in  11.)  acetic  acid  is  only  .004  ionized 
(p.  228),  so  that,  in  the  equation  for  the  equilibrium, 

(.996)  HCaH3Oa<=?H'-(.004)  +  C2H3(V  (.004), 

the  relative  proportions  are  as  shown  by  the  numbers  in  parenthesis. 
If  the  whole  of  the  acid  (60  g.)  were  ionized,  there  would  be  1  g.  of 
hydrogen-ion  per  liter.  Yet,  even  in  this  much  smaller  concen- 
tration (.004  g.  per  liter),  the  acid  taste  of  the  H*  and  its  effect  upon 
indicators  can  be  distinctly  recognized.  If,  now,  solid  sodium 
acetate  is  dissolved  in  the  solution,  the  liquid  no  longer  gives  an  acid 
reaction  with  one  of  the  less  delicate  indicators,  like  methyl  orange 
(q.v.).  The  explanation  is  simple.  Sodium  acetate  is  highly 
ionized.  It  gives,  therefore,  a  large  concentration  of  acetate-ion  to 
a  liquid  formerly  containing  very  little.  This  causes  a  greatly 
increased  union  of  the  H*  ions  and  C2H3O2'  ions  and  the  former, 
being  already  very  few  in  number,  disappear  almost  entirely. 
Hence  the  solution  becomes,  to  all  intents  and  purposes,  neutral. 
There  is  no  less  acetic  acid  present  than  before,  but  the  concentration 
of  hydrogen-ion  is  very  much  smaller. 

Formulation  and  Quantitative  Treatment  of  the  Case  of 
Excess  of  One  Ion.  —  If  the  semi-mathematical  mode  of  formu- 
lating an  equilibrium  (p.  180),  as  applied  to  the  case  of  an  ionogen 
(p.  224),  be  employed  here,  the  foregoing  general  statements  may  be 
made  more  precise  and  the  conclusions  clearer.  If  [H'j  and  [C2H302'] 
represent  the  molecular  concentrations  of  hydrogen-ion  and  acetate- 
ion,  respectively,  and  [HC2H302]  that  of  the  acetic  acid  molecules  at 
equilibrium,  then: 

[IT]  X  [C2H30/1  _  K 
[HC2H,OJ 

The  value  of  K  is  constant,  whether  the  strength  of  the  solution  of 
acetic  acid  is  great  or  small,  and  even  when  another  substance  with 
a  common  ion  is  present.  In  the  latter  case,  [C2H3O2']  and  [H*] 
stand  for  the  whole  concentrations  of  each  of  these  ionic  substances 
from  both  sources. 

Now,  in  normal  acetic  acid  [H*]  =  .004,  [C2H302/]  =  .004  (for  the 
number  of  each  kind  of  ions  is  the  same),  and  [HC2H3O2]  =  .996, 
practically  1.  Substituting  in  the  formula: 

0.004  x  0.004  =g(=0.Q4l6). 


IONIC    EQUILIBRIUM,    CONSIDERED    QUANTITATIVELY         383 

When  sodium  acetate  is  dissolved  in  the  liquid  until  the  solution  is 
normal  in  respect  to  this  substance  also,  the  following  additional 
equilibrium  has  to  be  considered: 

(.47)  NaC3H3O2  <=±  Na*  (.53)  +  C2H3O2'  (.53). 

The  concentration  of  acetate-ion  from  this  source  is  .53,  so  that,  m 
the  mixture  of  acid  and  salt,  the  concentration  of  acetate-ion 
[CaHgO/]  will  be  .53  +  .004  =  .534,  or  nearly  134  times  larger  than 
in  the  acid  alone.  Hence,  in  order  that  the  product  [H*]  X  [C2H3O2/] 
may  recover,  as  it  must,  a  value  much  nearer  to  the  old  one,  [H*] 
must  be  diminished  to  something  like  T^¥  of  its  former  magnitude. 
That  is,  [H*]  will  become  equal  to  about  0.00003,  the  rest  of  the 
hydrogen-ion  uniting  with  a  corresponding  amount  of  the  acetate-ion 
to  form  molecular  acetic  acid.  The  effect  of  adding  this  amount  of 
sodium  acetate  therefore  is,  as  we  have  seen,  to  reduce  the  concen- 
tration of  the  hydrogen-ion  below  the  amount  which  can  be  detected 
by  use  of  an  indicator  like  methyl-orange. 

This  effect  is  of  course  reciprocal,  and  the  ionization  of  the  sodium 
acetate  will  be  reduced  also.  But  the  acetate-ion  furnished  by  the 
acetic  acid  is  relatively  so  small  in  amount  (.00003  against  .53)  that 
the  effect  it  produces  on  the  ionization  of  the  salt  is  imperceptible. 

It  will  be  noted  that  the  acetate-ion  and  hydrogen-ion  disappear 
in  equivalent  quantities,  for  they  unite.  There  is,  however,  so  much 
of  the  former  that  its  loss  goes  unremarked,  while  there  is  so  little  of 
the  latter  that  almost  none  of  it  remains.  When  substances  of  more 
nearly  equal  degrees  of  ionization  are  used,  both  effects  are  equally 
inconspicuous.  Thus,  sodium  chloride  and  hydrogen  chloride  in 
normal  solutions  yield  approximately  equal  concentrations  of 
chloride-ion  (.784  and  .676).  Hence,  if  one  mole  of  sodium  chloride 
were  to  be  dissolved  in  the  portion  of  water  already  containing  one 
mole  of  hydrogen  chloride,  the  concentration  of  the  chloride-ion,  at 
a  very  rough  estimate,  would  be  nearly  doubled.  If  this  doubling 
of  the  concentration  of  chloride-ion  almost  halved  that  of  the 
hydrogen-ion  (.784),  in  order  that  the  expression  [Cl'J  X  [H']  -r  [HC1] 
might  remain  constant,  the  concentration  of  the  hydrogen-ion  would 
still  be  about  .400  and  therefore  100  times  as  great  as  in  molar  acetic 
acid.  It  is  thus  altogether  impossible  to  reduce  the  concentration 
of  the  hydrogen-ion  given  by  an  active  acid  like  hydrochloric  acid 
below  the  limit  at  which  indicators  are  affected,  for  there  is  no  way 


384  COLLEGE    CHEMISTRY 

of  introducing  the  enormous  concentration  of  the  other  ion  which 
the  theory  demands. 

With  more  crude  means  of  observation  than  indicators  afford, 
effects  like  this  last  may  sometimes  be  rendered  visible.  This  was 
the  case  with  cupric  bromide  solution,  to  which  potassium  bromide 
was  added  (p.  235).  The  blue  of  the  cupric-ion  disappeared  from 
view,  while  much  cupric-ion  was  still  present,  because  the  brown 
color  of  the  molecular  cupric  bromide  covered  it  up  completely. 

Special  Case  of  Saturated  Solutions.  —  The  commonest  as 
well  as  the  most  interesting  application  of  the  conceptions  developed 
above  is  met  with  in  connection  with  saturated  solutions,  especially 
those  of  relatively  insoluble  substances. 

The  situation  in  a  system  consisting  of  the  saturated  solution  and 
excess  of  the  solute  has  been  discussed  already  (p.  102).  In  the  case 
of  potassium  chlorate,  for  example,  we  have  the  following  scheme  of 
equilibria: 

KC103  (solid)  <=±  KC103  (dslvd)  <r±  K*  +  C103'. 

Solution  of  the  solid  is  promoted  by  the  solution  pressure  of  the  mole- 
cules, while  it  is  opposed  by  the  osmotic  pressure  of  the  dissolved 
substance,  and  the  solution  is  saturated  when  these  tendencies 
produce  equal  effects  (p.  103).  Now  it  must  be  noted  that  the 
tendency  directly  opposed  to  the  solution  pressure  is  the  partial 
osmotic  pressure  of  the  dissolved  molecules  alone.  The  chief  con- 
tents of  the  solution,  the  molecules  and  two  kinds  of  ions  of  the  salt, 
and  any  foreign  material  that  may  be  present,  are  like  a  mixture  of 
gases,  and  the  principle  of  partial  pressure  (p.  60)  is  to  be  applied. 
The  ions  and  the  foreign  material  do  not  deposit  themselves  upon  the 
solid,  and  take,  therefore,  no  part  directly  in  the  equilibrium  which 
controls  solubility.  In  respect  to  this  the  ions  are  themselves 
foreign  substances.  Hence  the  conclusion  may  be  stated  that,  in 
solutions  saturated  at  a  given  temperature  by  a  given  solute,  the  concen- 
tration of  the  dissolved  molecules  of  the  solute  considered  by  themselves 
will  be  constant  whatever  other  substances  may  be  present. 

The  total  solubility  of  a  substance,  as  we  have  used  the  term 
hitherto,  is  made  up  of  a  molecular  and  an  ionic  part.  The  latter, 
as  we  shall  presently  see,  is  not  constant  when  a  foreign  substance 
containing  a  common  ion  is  already  in  the  liquid.  Since  the  treat- 


IONIC    EQUILIBRIUM,    CONSIDERED   QUANTITATIVELY          385 


ment  of  the  subject  requires  us  now  to  distinguish  between  the  two 
portions  of  the  solute,  a  diagram  (Fig.  65)  will  assist  in  emphasizing 
the  distinction.  The  material  at  the  bottom  is  the  salt.  The 
molecules  and  ions  are  to  be  thought  of  as  being  mixed  and  as  being 
present  in  numbers  represented  by  the  factors  n  and  m.  Since  no 
foreign  body  is  present,  the  two  ions  in  this  case  are  equal  in  number. 
When  we  now  apply  these  ideas  to  the  mathematical  expression 
of  the  relation: 

[K1  X  [010,1 
[KG1OJ 

we  perceive  that,  in  a  saturated  solution,  [KC1O3],  the  concentration 
of  the  molecules,  is  constant.     Transposing,  we  have 

[K-]  X  [CIO/]  -  K  [KC10,']  =  K\ 

Hence  the  relation  leads  to  the  important  conclusion  that,  in  a 
saturated  solution  the  product  of  the  molar  concentrations  of  the  ions  is 
constant.*  This  product  is  called  the  ion-product 
constant  for  the  substance.  The  law  of  the  con- 
stancy of  the  ion-product  in  a  saturated  solution 
is  one  of  the  most  useful  of  the  principles  of 
chemistry.  It  enables  us  to  explain  all  the 
varied  phenomena  of  precipitation  and  of  the 
solution  of  precipitates  in  a  consistent  manner. 
These  applications  of  the  principle  will  be 
explained  in  the  next  chapter.  One  curious 
kind  of  precipitation  will  be  described  here, 
however,  as  an  illustration  of  the  use  of  the 
principle. 


mK'+mCIOs' 


n  KCIOs 


FIG.  65. 

Illustration    of   the    Principle    of   Ion- 
Product  Constancy.  —  When,  to  a  saturated  solution  of  one  of  the 
less  soluble  salts,  a  strong   solution   of  a  salt  having  one  ion  in 
common  with  the  first  salt  is  added,  precipitation  of  the  first  salt 
frequently    takes    place.      This    happens,    for    example,    with   a 

*  The  principle  of  constant  concentration  of  dissolved  molecules,  stated 
above,  has  been  shown  to  express  the  facts  very  inaccurately.  Now  the 
principle  of  the  constancy  of  the  ratio  of  the  ion-product  to  the  concentration 
of  the  molecules  is  also  inaccurate,  yet  in  such  a  way  that  the  two  errors 
neutralize  one  another.  Thus,  the  principle  of  ion-product  constancy  here 
given  is  in  itself  fairly  exact. 


386  COLLEGE   CHEMISTRY 

saturated  solution  of  potassium  chlorate,  which  is  not  very 
soluble  (molar  solubility  .52,  see  Table).  The  concentrations  [K*] 
and  [C103X]  being  small,  one  may  easily  increase  the  value  for 
one  of  the  ions,  say  [CIO/],  fivefold,  by  adding  a  chlorate  which 
is  sufficiently  soluble.  To  preserve  the  value  of  the  product 
[K']  X  [C103'],  the  value  of  [K*]  will  then  have  to  be  diminished 
at  once  to  one-fifth  of  its  former  value.  This  can  occur  only 
by  union  of  the  ionic  material  it  represents  with  an  equivalent 
amount  of  that  for  which  [CIO/]  stands.  The  molecular  material 
so  produced  will  thus  tend  at  first  to  swell  the  value  of  [KC103]. 
But  the  value  of  [KC1O3]  cannot  be  increased,  for  the  solution  is 
already  saturated  with  molecules,  so  that  the  new  supply  of  molecules, 
or  others  in  equal  numbers,  will  be  precipitated.  Hence  the  ionic 
part  of  the  dissolved  substance  may  be  diminished,  the  equilibria 
(p.  384)  may  be  partially  reversed,  and  we  may  actually  precipitate 
a  part  of  the  dissolved  material  without  introducing  any  substance, 
which,  in  the  ordinary  sense,  can  interact  with  it. 

In  point  of  fact,  when,  to  a  saturated  solution  of  potassium 
chlorate  there  is  added  a  saturated  solution  of  potassium  chloride 
(KC1)  or  of  sodium  chlorate  (NaClO3),  a  precipitate  of  potassium 
chlorate  is  thrown  down.  These  two  salts,  containing  each  one  of 
the  ions  of  KC1O3,  and  being  much  more  soluble  than  the  latter 
(see  Table),  increase  the  concentration  of  one  ion  and  cause  the 
precipitation  in  the  fashion  just  explained. 

The  product  of  the  concentrations  of  the  ions,  for  example  [K*] 
X  [CIO/],  is  called  also  the  solubility  product,  because  these  two  values 
jointly  determine  the  magnitude  of  the  solubility  of  the  substance. 
The  solubility  of  the  molecules  is  irreducible,  but  the  ionic  part  of 
the  dissolved  material  may  become  vanishingly  small  if  the  value 
of  either  [X*]  or  [Y'J  is  very  minute.  The  ionic  part  of  any  particular 
substance  is  made  up  of  the  smaller  of  the  two  concentrations,  of  the 
ionic  substances  which  it  yields,  plus  an  equivalent  amount,  and  no 
more,  of  the  concentration  of  the  other  ion.  The  rest  of  the  other 
ionic  substance  is  part  of  the  solubility  of  some  other  component. 

Other  Illustrations.  — -  The  precipitation  of  sodium  chloride 
from  a  saturated  solution,  by  the  introduction  of  gaseous  hydrogen 
chloride  (p.  374),  is  to  be  explained  in  the  same  manner.  The 
equilibria: 

NaCl  (solid)  «=±  NaCl  (dslvd)  ^±  Na"  +  Cl" 


IONIC    EQUILIBRIUM,    CONSIDERED    QUANTITATIVELY         387 

are  reversed  by  the  introduction  of  additional  Cl'  from  the  very 
soluble,  and  highly  ionized  HCL 

The  evolution  of  a  steady  stream  of  hydrogen  chloride  is  often 
accomplished  by  allowing  concentrated  sulphuric  acid  to  flow  into 
saturated  hydrochloric  acid: 

H'  +  Cl7  <=>  HC1  (dslvd)  <±  HC1  (gas). 

The  effect  is  due  in  part  to  repression  of  the  ionization  of  the  hydro- 
gen chloride  and  elimination  of  molecules  of  the  gas  from  the  water 
which  is  already  saturated  with  molecules  of  the  same  kind.  The 
"  salting  out  "  of  soap  (p.  335)  is  probably  a  similar  phenomenon. 

Exercises.  —  1.  The  vapor  density  of  sodium  peroxide  has  not 
been  determined.  Why  is  the  formula  Na2O2  assigned  to  it? 

2.  Construct  a  scheme  of  equilibria  (p.  253)  showing  the  hydroly- 
sis of  calcium  sulphide.    Why  does  the  presence  of  calcium  hydroxide 
diminish  the  tendency  to  hydrolysis? 

3.  Show  the  application  of  the  principle  of  ion-product  constancy 
to  the  salting  out  of  soap  (p.  335). 

4.  What  will  be  the  effect  of  adding  a  concentrated  solution  of 
silver  nitrate  to  a  saturated  solution  of  silver  sulphate  (see  Table  of 
solubilities)? 


CHAPTER  XXXV 
THE   METALLIC   ELEMENTS   OF    THE   ALKALINE   EARTHS 

The  Chemical  Relations  of  the  Elements.  —  The  metals 
of  this  group,  calcium  (Ca,  at.  wt.  40.1),  strontium  (Sr,  at.  wt.  87.6), 
and  barium  (Ba,  at.  wt.  137.4),  constitute  a  typical  chemical  family, 
both  in  the  qualitative  resemblance  to  one  another  of  the  elements 
and  of  the  corresponding  compounds,  and  in  the  quantitative  varia- 
tion in  the  properties  with  increasing  atomic  weight.  The  metals 
themselves  displace  hydrogen  vigorously  from  cold  water,  giving 
hydroxides.  The  solutions  of  these  hydroxides,  although  dilute, 
on  account  of  a  rather  small  solubility,  are  strongly  alkaline  in 
reaction.  The  high  degree  of  ionization  of  the  hydroxides  recalls 
the  hydroxides  of  the  metals  of  the  alkalies,  and  their  relative 
insolubility  the  hydroxides  of  the  "  earths  "  (q.v.). 

In  all  their  compounds,  calcium,  strontium,  and  barium  are  biva- 
lent. The  hydroxides  are  formed  by  union  of  the  oxides  with  water, 
and  are  progressively  less  easy  to  decompose  by  heating,  barium  hy- 
droxide being  the  hardest.  The  carbonates,  when  heated,  yield  the 
oxide  of  the  metal  and  carbon  dioxide,  barium  carbonate  being  the 
most  difficult  to  decompose.  The  nitrates,  when  heated  moderately, 
give  the  nitrites,  but  the  latter  are  broken  up  by  further  heating 
and  yield  the  oxide  of  the  metal,  and  nitrogen  tetroxide.  In 
these  and  other  respects  the  compounds  of  the  metals  of  the  alkaline 
earths  resemble  those  of  the  heavy  metals  and  differ  from  those  of 
the  metals  of  the  alkalies.  Barium  approaches  the  latter  most 
nearly. 

The  table  of  solubilities  (q.v.)  shows  that  the  chlorides  and  nitrates 
of  calcium,  strontium,  and  barium  are  all  soluble  in  water,  the 
solubility  diminishing  in  the  order  given.  The  sulphates  and 
hydroxides  cover  a  wide  range  from  slight  solubility  to  extreme 
insolubility.  Of  the  sulphates,  2100, 1 10,  and  2.3  parts,  respectively, 
dissolve  in  one  million  parts  of  water.  In  the  case  of  the  hydroxides 
the  order  of  magnitude  is  reversed,  and  the  corresponding  numbers 

388 


CALCIUM  389 

are  200,  630,  and  2200.  The  carbonates  are  almost  as  insoluble  as 
is  barium  sulphate.  The  new  element,  radium  (Ra,  at.  wt.  226.5), 
appears  to  belong  to  this  family  (see  under  Uranium). 

CALCIUM  Ca. 

Occurrence.  —  The  fluoride,  and  the  various  forms  of  the  car- 
bonate, sulphate,  and  phosphate  which  are  found  in  nature,  are 
described  below.  As  silicate,  calcium  occurs,  along  with  other 
metals,  in  many  minerals  and  rocks.  The  element  is  found  also  in 
plants,  and  its  compounds  are  constituents  of  the  bones  and  shells  of 
animals. 

The  Metal.  —  Calcium  is  made  by  electrolysis  of  the  molten 
chloride.  A  hollow  cylinder  made  of  blocks  of  carbon  bolted 
together  and  open  above,  forms  the  anode.  A  rod  of  colfer  hang- 
ing so  that  its  end  dips  into  the  melt  forms  the  cathode,  xne  melt- 
ing of  the  anhydrous  calcium  chloride  with  which  the  cylinder  is 
filled  is  started  by  means  of  a  thin  rod  of  carbon  laid  across  from  the 
anode  to  the  cathode.  When  the  heat  generated  by  the  passage  of 
the  current  through  this  highly  resisting  medium  has  melted  a 
sufficient  amount  of  the  salt,  the  rod  is  removed,  and  the  resistance 
of  the  fused  material  suffices  to  maintain  the  temperature.  The 
calcium  rises  round  the  cathode  and  collects  on  the  surface  of  the 
bath.  By  slowly  elevating  the  copper  cathode,  the  calcium,  which 
adheres  to  it,  may  be  drawn  out  of  the  fused  mass  in  the  form  of  a 
gradually  lengthening,  irregular  rod.  The  rod  of  calcium  is  kept 
constantly  in  contact  with  the  metal  which  accumulates  on  the 
surface,  and  thus  forms  one  of  the  electrodes. 

Calcium  is  a  silver-white,  crystalline  metal  (m.-p.  760°,  sp.  gr.  1.85) 
which  is  a  little  harder  than  lead,  and  can  be  cut,  drawn,  and  rolled. 
It  interacts  rapidly  with  water.  When  dry  and  cold  it  is  inactive, 
but  when  heated  it  unites  vigorously  with  hydrogen,  oxygen,  the 
halogens,  and  nitrogen.  On  this  account  it  is  used  in  producing 
a  high  degree  of  evacuation.  It  burns  in  the  air,  giving  a  mixture 
of  the  oxide  and  nitride  Ca3N2.  The  presence  of  the  latter  may  be 
shown  by  the  liberation  of  ammonia  when  water  is  added  to  the 

reSldue:  Ca,N2  +  6H20  -»  3Ca(OH)2  +  2NH8. 


390  COLLEGE   CHEMISTRY 

A  white  crystalline  hydride  CaH2  is  formed  by  direct  union  of 
the  constituents.  It  is  known  in  commerce  as  hydrolyte. 

The  manufacture  of  calcium  carbide  CaC2  (p.  321),  and  the  forma- 
tion of  acetylene  by  its  interaction  with  water  (p.  330),  have  already 
been  described. 

Calcium  Chloride  CaCl2* —  This  salt,  for  which  there  is  no 
extensive  commercial  application,  is  formed  as  a  by-product  in 
several  industrial  operations.  Thus,  it  arises  in  the  liberation  of 
ammonia  from  ammonium  chloride  by  the  action  of  lime,  in  the 
manufacture  of  potassium  chlorate  (p.  195),  and  in  the  Solvay  soda 
process  (p.  378).  By  evaporation  of  any  solution,  the  hexahydrate 
CaCl2,6H2O  is  obtained  in  large,  deliquescent  ,^^-sided  prisms.  On 
account  of  the  great  concentration  of  a  satiSRed  solution  of  this 
compound,  the  solid  and  solution  do  not  reach  a  condition  of  equi- 
librium j^h  ice  (cf.  p.  204)  until  the  temperature  has  fallen  to 
—  48°.  ^Rie  hydrate,  when  mixed  with  ice,  gives,  therefore,  a  very 
efficient  freezing  mixture. 

Calcium  chloride,  dehydrated  by  heating  (CaCl2),  forms  a  porous 
mass  which  is  used  in  chemical  laboratories  for  drying  gases  and 
liquids. 

Calcium  chloride  forms  compounds,  not  only  with  water,  but  also 
with  ammonia  (CaCl2,8NH3)  and  with  alcohol.  For  drying  these 
substances,  therefore,  quicklime  is  employed. 

Calcium  Fluoride  CaF2.  —  This  compound  occurs  in  nature  as 
fluorite  or  fluor-spar  CaF2.  It  crystallizes  in  cubes,  is  insoluble  in 
Water,  and  when  pure  is  colorless.  Natural  specimens  often  possess 
a  green  tint  or  show  a  violet  fluorescence.  It  is  formed  as  a  precipi- 
tate when  a  soluble  fluoride  is  added  to  a  solution  of  a  salt  of  calcium. 

Fluorite  is  used  in  the  etching  of  glass,  as  the  source  of  the  hydro- 
gen fluoride  (p.  171).  It  is  easily  fusible,  as  its  name  indicates  (Lat. 
fluere,  to  flow),  and  is  employed  in  metallurgical  operations  as  a 
flux  (p.  355),  for  lowering  the  melting-point  (or  freezing-point, 
which  is  the  same  thing,  cf.  p.  204)  of  the  slag  (p.  355),  and  so 
facilitating  the  separation  of  the  latter  from  the  metal. 

Calcium  Carbonate  CaCO3.  —  This  compound  is  found  very 
plentifully  in  nature.  Limestone  is  a  compact,  indistinctly  crystal- 


CALCIUM  391 

line  variety,  while  marble  is  a  distinctly  crystalline  form.  Chalk  * 
is  a  deposit  consisting  of  the  calcareous  parts  of  minute  organisms. 
Egg-shells,  oyster-shells,  coral,  and  pearls  are  other  varieties  of 
organic  origin,  f  Calcite  and  Iceland  spar  (Ger.  spalten,  to  split) 
are  pure  crystallized  calcium  carbonate.  The  former  occurs  in  flat 
rhombohedrons,  or  in  pointed,  six-sided  crystals  (Fig.  33,  p.  94) 
known  as  scalenohedrons  ("  dog-tooth  "  spar)  belonging  to  the  same 
system. 

When  heated,  calcium  carbonate  dissociates,  giving  carbon  dioxide 
and  quicklime:  CaCO3  ^  CaO  +  CO, 

At  ordinary  temperatures  the  decomposition  is  imperceptible.  On 
the  contrary,  atmospheric  carbon  dioxide,  in  spite  of  its  very  low 
partial  pressure,  combines  with  quicklime,  giving  "  air-slaked  "  lime. 
As  the  temperature  rises,  however,  the  tension  of  carbon  dioxide  coming 
from  the  carbonate  increases,  and  has  a  fixed  value  for  each  temperature. 
If  it  is  continually  allowed  to  escape,  so  that  the  maximum' pressure 
is  not  reached,  the  whole  of  the  salt  eventually  decomposes.  At 
812°  the  pressure  reaches  one  atmosphere,  and  at  865°  two  atmos- 
pheres. The  phenomenon  is  precisely  similar  to  the  dissociation  of 
barium  dioxide  (cf.  Erin's  process,  pp.  46,  184)  and  to  the  evapor- 
ation of  a  liquid  (pp.  78-79) . 

Limestone  is  used  in  the  manufacture  of  quicklime  (q.v.)  and  of 
glass.  It  is  employed  largely  as  a  flux  in  metallurgy,  when  min- 
erals rich  in  silica  are  brought  into  fusible  form  by  the  production 
of  calcium  silicate  (CaSi03).  Large  amounts  also  find  application 
as  building-stone. 

Hard  Water. —  Calcium  carbonate  is  almost  insoluble  in  water. 
In  the  cold  the  solubility  is  a  little  over  1  part  in  100,000;  in  hot 
water  the  solubility  is  even  smaller.  Water  containing  carbonic 
acid  dissolves  it  more  freely,  on  account  of  the  formation  of  the  more 
soluble  calcium  bicarbonate  (cf.  p.  324) : 

CaC03  +  H2C03  <±  Ca(HC03)2. 

At  15°  a  liter  of  water  saturated  with  carbon  dioxide  (at  760  mm. 
pressure)  dissolves  0.385  g.  of  the  carbonate,  whereas  a  liter  of  pure 

*  Blackboard  "crayon"  is  usually  made  of  gypsum  and  not  of  chalk, 
t  The  hard  coverings  of  Crustacea  and  insects  are  not  made  of  this  sub- 
stance, but  of  an  organic  material  called  chitin. 


392  COLLEGE   CHEMISTRY 

water  dissolves  but  .013  g.  With  a  higher  pressure  of  carbon  dioxide 
still  larger  amounts  may  be  dissolved.  Conversely,  when  the 
carbon  dioxide  is  driven  out  by  boiling,  the  carbonate  is  repre- 
cipitated. 

Water  containing  salts  of  lime  or  magnesia  in  solution  is  known 
as  hard  water.  The  carbonate,  which  may  be  thrown  down  by 
boiling,  gives  "temporary  hardness,"  while  the  sulphate  of  calcium, 
since  it  is  soluble  per  se  and  is  not  affected  by  boiling,  gives  "  per- 
manent hardness  "  (for  the  action  on  soap,  see  p.  335). 

The  temporary  hardness  may  be  removed  by  adding  to  the  water 
a  quantity  of  slaked  lime  sufficient  to  convert  the  excess  of  carbonic 
acid  into  calcium  carbonate.  The  permanent  hardness  may  be 
destroyed  by  the  addition  of  sodium  carbonate.  In  both  cases  the 
precipitated  carbonate  is  allowed  to  settle  or  is  separated  by  filtra- 
tion. 

When  evaporated  in  steam-boilers  hard  water  leaves  behind  a 
heavy  deposit  of  boiler-crust  containing  all  the  salts  formerly  in 
solution.  Subterranean  water,  rich  in  carbon  dioxide,  finding  its 
way  to  the  roofs  of  caverns,  loses  its  gas  and  gives  rise  to  stalactites. 
The  drippings  form  stalagmites  on  the  floors. 

Calcium  Oxide  and  Hydroxide.  —  Pure  oxide  of  calcium  CaO 
(quicklime)  may  be  made  by  ignition  of  pure  marble  or  calcite.  For 
commercial  purposes  limestone  is  converted  into  quicklime  in  kilns. 
In  the  United  States  the  "  long-flame  "  process,  in  which  the  kiln  is 
first  charged  with  limestone  and  a  fire  is  then  kindled  in  a  cavity 
left  at  the  bottom,  is  the  one  most  commonly  used.  Often,  the  lime- 
stone acfi  coal  are  simply  thrown,  in  alternate  layers,  into  the  kiln,  and 
the  products  are  withdrawn  at  the  bottom.  The  latter,  the  "  short- 
flame  "  method,  demands  less  fuel,  since  the  operation  is  continuous, 
and  the  structure  is  never  allowed  to  cool,  but  the  quicklime  is 
mixed  with  the  ash  of  the  coal.  Ih  many  cases  this  serious  objec- 
tion is  removed  by  placing  the  fuel  in  cavities  at  the  sides  of  the 
kiln. 

Pure  calcium  oxide  is  a  white,  porous  solid.  It  is  barely  fusible 
in  the  oxyhydrogen  flame,  but  may  be  melted  and  boiled  in  the 
electric  arc.  It  is  not  reducible  by  sodium,  or  by  carbon  excepting 
at  the  temperature  of  the  electric  furnace. 

When  water  is  poured  upon  quicklime,  it  is  first  absorbed  into  the 


CALCIUM  393 

pores  mechanically,  and  then  unites  chemically  to  form  calcium 
(yrdroxid@(CaOH)2: 

CaO  +  H2Ot=>Ca(OH)2. 

The  product  is  a  bulky  powder.  The  actfc@is  accompa^A  by  the 
developing  of  much  heat.  The  charge  is  ^ersible,  anc^it  a  high 
temp^jature  the  hydro^e  can  be  dehydrated. 

Calcium  hydroxide  is  (gjightly  soluble  in  \\gfcer,  —  1  part  in  600 
parts  of  water  at  18°,  about  twice  as(5Q^)w^r  being  required  £ 
100°.  The  solution,  relatively  to  its  concentration,  is  strongly 
alkaline.  On  account  of  its  cheapness,  this  substance  is  used  by 
manufacturers  in  almost  all  operations  requiring  a@ise,  and  it  thfls 
occupi^tt©same  position  amongst  bases  that  sulphuric  acj^does 
amongst  acids.  Some  of  the  industries  in  which  caustic  lime  plays 
a  part  are:  the  manufacture  of  alkalies  (p.  363),  bleaching  powder, 
and  mortar  (see  below),  the  removal  of  the  hair  from  hides  in  prep- 
aration for  tanning,  and  the  purification  of  illuminating-gas  (p.  340). 

Mortar  and  Cement.  —  Mortar  is  made  by  mixing  water  with 
slaked  lime  and  a  large  proportion  of  sand.  The  "  hardening  " 
process  consists  in  an  interaction  of  the  carbon  dioxide  of  the  air 
with  the  calcium  hydroxide: 

C02  +  Ca(OH)2  ->  CaC03  +  H20. 

After  the  superficial  parts  have  been  changed,  the  process  goes  on 
very  slowly,  and  many  years  are  required  before  the  deeper  layers 
have  been  transformed.  The  minute  crystals  of  calcium  carbonate 
which  are  formed  are  interlaced  with  the  sand  particles,  and  a  rigid 
mass  is  finally  produced.  The  "  hardening  "  does  not^^gin  until 
the  excess  of  water  used  in  making  the  n©tar  has  evaporated,  and 
hence  ordinary  mortar  is  unsuitable  for  use  in  damp  places  such  as 
cellars. 

Cement  is  made  by  strong^)  heating  a  mixture  of  limestone,  clay, 
and  sand,  and  pulverizing  the  product.  When  the  cement  is  mixed 
with  water,  it  gradually  sets  to  a  solid  mass  which  appears  to  consist 
of  a  mixture  of  silicates  of  calcium  and  aluminium.  Since  the 
change  is  not  dependent  on  access  of  any  gas,  it  proceeds  through- 
out the  whole  material  simultaneously,  and  not,  as  in  the  case  of 
mortar  from  the  surface  inwards.  For  this  reason  the  hardening  of 
cement  occurs  just  as  well  under  water  as  in  any  other  locality. 


394  COLLEGE   CHEMISTRY 

Calcium  Oxalate  Ca€2O4.  —  This  salt  may  be  observed  under 
the  microscope  in  the  cells  of  many  plants.  It  appears  in  the  form 
of  needle-shaped  or  of  granular  crystals.  Since  it  is  the  least  soluble 
salt  of  calcium,  its  formation  by  precipitation  is  used  as  a  test  for 
calcium  ions. 

Theory  of  Precipitation.  —  The  precipitation  of  calcium  oxa- 
late CaC2O4,  just  referred  to,  is  a  typical  one  and  may  be  used  to 
illustrate  the  application  of  ion-product  constancy  (p.  385)  to 
explaining  the  phenomenon.  The  same  explanation  serves  for  all 
precipitations. 

The  first  thing  to  be  remembered  is  that  the  precipitate  which  we 
observe,  however  insoluble  its  material  may  be,  does  not  include  oil 
of  the  substance,  but  only  the  excess  beyond  what  is  required  to 
saturate  the  water.  The  liquid  surrounding  the  precipitate  is  always  a 
saturated  solution  of  the  substance  precipitated.  If  it  were  not  so, 
some  of  the  precipitate  would  dissolve  until  the  liquid  became 
saturated.  Thus,  for  example,  when  we  add  ammonium  oxalate 
solution  to  calcium  chloride  solution: 


the  liquid  is  a  saturated  solution  of  calcium  oxalate,  with  the  excess 
of  this  salt  suspended  in  it. 

Looking  at  the  matter  from  this  view-point,  we  perceive  the  appli- 
cation of  the  rule  of  ion-product  constancy.  In  this  saturated  solu- 
tion (p.  384)  the  product  of  the  ion-concentrations,  [Ca**]  X  [C2O4"], 
is  constant.  If  the  original  solutions  had  been  so  very  dilute  that, 
when  they  were  mixed,  the  product  of  the  concentrations  of  these 
two  ions  had  not  reached  the  value  of  this  constant,  no  precipitation 
would  have  occurred.  As  a  matter  of  fact  the  ion-product  considerably 
exceeded  the  requisite  value,  and  hence  the  salt  was  thrown  down 
until  the  balance  remaining  gave  the  value  in  question.  The  rule 
for  precipitation,  then,  is  as  follows:  Whenever  the  product  of  the  con- 
centrations of  any  two  ions  in  a  mixture  exceeds  the  value  of  the  ion- 
product  in  a  saturated  solution  of  the  compound  formed  by  their  union, 
this  compound  will  be  precipitated.  Naturally  the  substances  with 
small  solubilities,  and  therefore  small  ion-product  constants,  are  the 
ones  most  frequently  formed  as  precipitates. 


THEORY  OF  PRECIPITATION  395 

Rule  for  Solution  of  Substances.  —  The  rule  for  solution  of 

any  ionogen  follows  at  once  from  the  foregoing  considerations,  and 
may  be  formulated  by  changing  a  few  of  the  words  in  the  rule  just 
given:  Whenever  the  product  of  the  concentrations  of  any  two  ions  in  a 
mixture  is  less  than  the  value  of  the  ion-product  in  a  saturated  solution 
of  the  compound  formed  by  their  union,  this  compound,  if  present  in  the 
solid  form,  will  be  dissolved.  When  applied  to  the  simplest  case, 
this  rule  means  that  a  substance  will  dissolve  in  a  liquid  not  yet 
saturated  with  it,  but  will  not  dissolve  in  a  liquid  already  saturated 
with  the  same  material.  The  value  of  the  rule  lies  in  its  application 
to  the  less  simple,  but  equally  common  cases,  such  as  when  an 
insoluble  body  is  dissolved  by  interaction  with  another  substance 
(next  section). 

Applications  of  the  Mule  for  Solution  to  the  Solution  of 
Insoluble  Substances.  —  So  long  as  a  substance  remains  in  pure 
water  its  solubility  is  fixed.  Thus,  with  calcium  hydroxide,  the 
system  comes  to  rest  when  .17  g.  per  100  c.c.  of  water  have  gone  into 
solution: 

Ca(OH)2  (solid)  <->  Ca(OH)2  (dslvd)  <=»Ca"  +  20H'. 

But  if  an  additional  reagent  which  can  combine  with  either  one  of 
the  ions  is  added,  the  concentration  of  this  ion  at  once  becomes  less, 
the  ion-product  therefore  tends  to  diminish,  and  further  solution 
must  take  place  to  restore  its  value.  Thus,  if  a  little  of  an  acid 
(giving  H*)  be  added  to  the  solution  of  calcium  hydroxide,  the  union 
of  OH'  and  H"  to  form  water  removes  the  OH',  and  solution  of  the 
hydroxide  proceeds  until  the  acid  is  used  up.  There  are  now  more 
Ca"  than  OH7  ions  present,  but  the  ion-product  reaches  the  same 
value  as  before,  and  then  the  change  ceases.  If  a  further  supply  of 
acid  is  added,  the  removal  of  OH'  to  form  H20  begins  again.  With 
excess  of  the  acid,  the  only  stable  OH'  concentration  is  that  which  is 
a  factor  in  the  very  minute  ion-product  of  water,  [OH']  X  [H*]. 
Hence,  the  calcium  hydroxide,  which  requires  in  general  a  much 
higher  concentration  of  OH'  than  this  to  precipitate  it  or  to  keep  it 
out  of  solution,  finally  all  dissolves. 

This  particular  action  is  a  neutralization  of  an  insoluble  base. 
But  the  other  kinds  of  actions  by  which  insoluble  ionogens  pass  into 
solution  all  resemble  it  closely,  and  differ  only  in  details.  The 
general  outlines  of  the  explanation  are  the  same  in  every  case.  We 


396  COLLEGE   CHEMISTRY 

proceed  now  to  apply  it  to  the  common  phenomenon  of  the  solution 
of  an  insoluble  salt  by  an  acid. 

Interaction  of  Insoluble  Salts  with  Acids,  Resulting  in 
Solution  of  the  Salt.  —  Calcium  oxalate  passes  into  solution 
when  in  contact  with  acids,  especially  active  acids.  Thus,  with 
hydrochloric  acid,  it  gives  calcium  chloride  and  oxalic  acid,  both  of 
which  are  soluble: 

CaC204 1  +  2HC1  <=±  CaCl2  +  H2C2O4.  (1) 

The  action  of  acids  upon  insoluble  salts  is  so  frequently  mentioned  in 
chemistry  and  is  so  important  a  factor  in  analytical  operations  that  it 
demands  separate  discussion.  This  example  is  a  typical  one  and 
may  be  used  as  an  illustration. 

According  to  the  rules  already  explained  (p.  394),  calcium  oxalate 
(or  any  other  salt)  is  precipitated  when  the  product  of  the  concen- 
trations of  the  two  requisite  ions  [Ca"]  X  [C2O4"]  exceeds  the  ion- 
product  for  a  saturated  solution  of  calcium  oxalate  in  pure  water. 
When,  on  the  contrary,  the  product  of  the  concentrations  of  the  two 
ions  falls  below  the  limiting  value,  a  condition  which  may  arise  from 
the  removal  in  some  way  either  of  the  Ca"  or  of  the  C204"  ions,  the 
undissociated  molecules  will  ionize,  and  the  solid  will  dissolve  to 
replace  them  until  the  ionic  concentrations  necessary  for  equilibrium 
with  the  molecules  have  been  restored  or  until  the  whole  of  the  solid 
present  is  consumed.  Here  the  oxalate-ion  from  the  calcium  oxa- 
late combines  with  the  hydrogen-ion  of  the  acid  (usually  an  active 
one)  which  has  been  added,  and  forms  molecular  oxalic  acid: 

C204"  +  2H*  i=»  H2C2O4.  (2) 

Hence,  dissociation  of  the  dissolved  molecules  of  calcium  oxalate  pro- 
ceeds, being  no  longer  balanced  by  encounters  and  unions  of  the  now 
depleted  ions,  and  this  dissociation  in  turn  leads  to  solution  of  other 
molecules  from  the  precipitate. 

It  will  be  seen  that  the  removal  of  the  ions  in  this  fashion  can 
result  in  considerable  solution  of  the  salt  only  when  the  acid  pro- 
duced is  a  feebly  ionized  one.  Here,  to  be  specific,  the  concentration 
of  the  C2O4"  in  the  oxalic  acid  equilibrium  (2)  above  must  be  less  than 
that  of  the  same  ion  in  a  saturated  calcium  salt  solution.  Now 
oxalic  acid  does  not  belong  to  the  least  active  class  of  acids,  and  its 
pure  solution  contains  a  considerable  concentration  of  C304". 


THEORY    OF   PRECIPITATION  397 

There  is,  lipwever,  a  decisive  factor  in  the  situation  which  we  have 
not  yet  taken  into  account.  The  hydrocHj)b&5  acid  which  we  used 
for  dissolving  th©  precipitate  is  a  vefy  highly  ionized  acid  and  gives 
an  enormously  greater  concentration  of  (hydrogen-ion  tl@i  does 
oxalic  acid.  Hence  the  l^drog£h-ion  is  in  exc$s  in  equation  (2), 

and  the  condition  of  equilibrium  for  oxalic  acid,  1?  J   X  K^O/'^  ^ 

[H2C2OJ 

will  be  satisfied  by  a  correspondingly  small  concentration  of  C204". 
In  this  particular  case,  therefore,  the  [C2O4"]  of  the  oxalic  acid  is 
less  than  that  given  by  the  calcium  oxalate.  The  whole  change, 
therefore,  depends  for  its  accomplishment  not  only  on  the  mere 
presence  of  hydrogen-ion,  but  on  the  repression  of  the  ionization  of  the 
oxalic  acid  by  the  great  excess  of  hydrogen-ion  furnished  by  the 
active  acid  that  has  been  used.  As  a  matter  of  fact,  we  find  that  a 
weak  acid  like  acetic  acid  has  scarcely  any  effect  upon  a  precipitate 
of  calcium  oxalate.  An  acid  stronger  than  oxalic  acid  must  be 
employed.  The  whole  scheme  of  the  equilibria  is  as  follows: 

CaC204(soUd)^CaC204(dsl vd)  *=,  Ca"  +  C2O4"l  _» H  c  O  (dslvd^  (^ 
2HC1  fc?2Cl/+  2H'  \*±°+4P&* 

When  excess  of  an  acid  sufficiently  active  to  furnish  a  large  concen- 
tration of  hydrogen-ion  is  employed,  the  last  equilibrium  is  then 
driven  forward  and  the  others  follow.  With  addition  of  a  weak  acid, 
only  a  slight  displacement  occurs,  and  the  system  comes  to  rest  again 
when  the  molecular  oxalic  acid  has  reached  a  sufficient  concentration. 

A  generalization  may  be  based  on  these  considerations:  an  insoluble 
salt  of  a  given  acid  will  in  general  interact  and  dissolve  when  treated 
with  a  solution  containing  another  acid,  provided  that  the  latter  acid  is  a 
much  more  highly  ionized  (more  active)  one  than  the  former  (see  below). 

But  even  active  acids  frequently  fail  to  bring  salts  of  weak  acids 
into  solution,  especially  when  the  weak  acid  is  itself  present  also. 
Here  the  cause  lies  in  the  fact  that  such  salts  are  less  soluble  than 
those  of  the  calcium  oxalate  type,  and  give  so  low  a  concentration 
of  the  negative  ion  that  the  utmost  repression  of  the  ionization  of 
the  corresponding  acid  does  not  give  a  lower  value  for  the  concen- 
tration of  this  ion  than  does  the  salt  itself.  Thus,  we  have  seen 
(p.  254)  that  even  hydrochloric  acid  (dilute)  will  not  dissolve  a 
number  of  sulphides.  For  example,  in  the  case  of  cupric  sulphide  in 
a  solution  saturated  with  hydrogen  sulphide,  the  S"  factor  in  the 


398  COLLEGE    CHEMISTRY 

solubility  product  [Cu"]  x  [S"]  remains  smaller  than  that  in  the 
scheme  defining  the  hydrogen  sulphide  equilibrium  [H']2  X  [S"] 
even  when  the  S"  factor  in  the  latter  is  diminished  in  consequence 
of  great  addition  of  hydrogen-ion.  In  this  case  the  first  link  in  the 
chain  of  equilibria: 

CuS  (solid)  fc5  CuS  (dsl vd)  ^Cu    +  S") 

2  HC1  *=>  2C1'  +  2H- }  ^  H2S     9lvd)' 

tends  so  decidedly  backward  that  only  the  use  of  concentrated  acid 
will  increase  the  concentration  of  the  H*  to  an  extent  sufficient  to 
secure  any  marked  advance  of  the  whole  action.  We  must  add, 
therefore,  to  the  above  rule:  provided  also  that  the  salt  is  not  one  of 
extreme  insolubility.  This  point  will  be  illustrated  more  fully  in 
connection  with  the  description  of  individual  sulphides  (see  under 
Cadmium). 

Illustrations  of  the  application  of  these  generalizations  are  count- 
less. Carbonic  acid  is  made  from  marble  (p.  322) ,  hydrogen  sulphide 
from  ferrous  sulphide  (p.  251),  hydrogen  peroxide  from  sodium 
peroxide  (p.  211),  and  phosphoric  acid  from  calcium  phosphate 
(p.  303).  In  each  case  the  acid  employed  to  decompose  the  salt  is 
more  active  than  the  acid  to  be  liberated.  On  the  other  hand,  cal- 
cium oxalate  is  insoluble  in  acetic  acid  because  this  acid  is  weaker 
than  is  oxalic  acid.  We  have  thus  only  to  examine  the  list  of  acids 
showing  their  degrees  of  ionization  (p.  228)  in  order  to  be  able  to 
tell  which  salts,  if  insoluble  in  water,  will  be  dissolved  by  acids,  and, 
in  general,  what  acids  will  be  sufficiently  active  in  each  case  for  the 
purpose.  In  chemical  analysis  we  discriminate  between  salts 
soluble  in  water,  those  soluble  in  acetic  acid  (the  insoluble  carbonates 
and  some  sulphides,  FeS  and  ZnS,  for  example),  those  requiring 
active  mineral  acids  for  their  solution  (calcium  oxalate  and  the  more 
insoluble  sulphides,  for  example),  and  those  insoluble  in  all  acids 
(barium  sulphate  and  other  insoluble  salts  of  active, acids). 

Precipitation   of  Insoluble   Salts  in  Presence  of  Acids. — 

The  converse  of  solution,  namely,  precipitation,  depends  upon  the  same 
conditions:  an  insoluble  salt  which  is  dissolved  by  a  given  acid  cannot 
be  formed  by  precipitation  in  the  presence  of  this  acid.  Thus,  calcium 
oxalate  can  be  precipitated  in  presence  of  acetic  acid,  but  not  in 
presence  of  active  mineral  acids  in  ordinary  concentrations.  Cupric 


CALCIUM  399 

sulphide  or  barium  sulphate  can  be  precipitated  in  presence  of  any 
acid,  but  ferrous  sulphide  and  jcalcium  carbonate  only  in  the  absence 
of  acids. 

From  this  it  does  not  follow  that  calcium  oxalate,  for  example, 
cannot  be  precipitated  if  once  an  active  acid  has  been  added  to  the 
mixture.  To  secure  precipitation,  all  that  is  necessary  is  to  remove 
the  excess  of  hydrogen-ion  which  is  repressing  the  ionization  of  the 
oxalic  acid.  This  can  be  done  by  adding  a  base,  which  removes  the 
H*,  or  even  by  adding  sodium  acetate.  The  acetate-ion  C2H3O2' 
unites  with  the  H*  to  form  the  little  ionized  acetic  acid,  in  presence 
of  which  calcium  oxalate  can  be  precipitated. 

Bleaching  Powder  Ca(OCl)Cl.  —  This  substance  (cf.  p.  189) 
is  manufactured  by  conducting  chlorine  into  a  box-like  structure 
containing  slaked  lime  spread  upon  perforated  shelves.  When  the 
transformation  has  reached  the  limit  (it  is  never  complete),  the 
supply  of  the  gas  is  shut  off,  and  some  lime-dust  is  blown  into  the 
chamber  to  absorb  the  remainder  of  the  free  chlorine. 

That  bleaching  powder  is  a  mixed  salt  CaCl(OCl)  rather  than  an 
equi-molar  mixture  of  calcium  chloride  and  calcium  hypochlorite, 
which  would  have  the  same  composition,  CaCl2,Ca(OCl)2,  is  rendered 
probable  by  the  facts  that  the  material  is  not  deliquescent  as  is 
calcium  chloride,  and  that  calcium  chloride  cannot  be  dissolved  out 
of  it  by  alcohol. 

Bleaching  powder  is  somewhat  soluble  in  water,  and  in  solution 
the  ions  Ca",  Cl',  and  CIO'  are  all  present.  Addition  of  active  acids 
causes  the  formation  of  hydrochloric  and  hypochlorous  acids  (p.  193). 
Weak  acids  like  carbonic  acid  displace  the  hypochlorous  acid  only 
(cf.  p.  191),  and  hence  the  dry  powder,  when  exposed  to  the  air,  has 
the  odor  of  the  latter  substance  rather  than  that  of  chlorine. 

The  substance  is  largely  used  by  bleachers  (cf.  p.  193),  and  as  a 
disinfectant  to  destroy  germs  of  putrefaction  and  disease. 

Calcium  Sulphate.— This  salt  is  found  in  large  quantities  in 
nature.  The  mineral  anhydrite  CaS04  occurs  in  the  salt  layers.  It 
contains  no  water  of  crystallization,  and  its  crystals  belong  to  the 
rhombic  system.  The  dihydrate,  CaSO4,2H20,  is  more  plentiful.  In 
granular  masses  it  constitutes  alabaster.  When  perfectly  crystal- 
lized (monoclinic  system,  Fig.  35,  p.  95)  it  is  named  gypsum  or 


400  COLLEGK  CHEMISTRY 

selenite.    The  same  hydrate  is  formed  by  precipitation  from  solutions. 
Its  solubility  is  about  1  in  500  at  18°. 

Plaster  of  paris  is  manufactured  by  heating  gypsum  until  nearly 
all  the  water  of  hydration  has  been  driven  out.  When  it  is  mixed 
with  water,  the  dihydrate  is  quickly  re-formed  and  a  rigid  mass  is 
produced.  That  the  plaster  sets  rapidly,  is  due  to  the  fact  that  the 
anhydrous  salt  is  more  soluble  than  the  dihydrate,  and  so  a  constant 
solution  of  the  one  and  deposition  of  the  other  goes  on  until  the  hy- 
dration is  complete.  It  becomes  rigid,  instead  of  forming  a  loose 
mass  of  dihydrate,  because  the  process  results  in  the  formation  of 
an  interlaced  and  coherent  mass  of  minute  crystals. 

CaS04  (solid)  *  CaSO,  (dslvd)  j  ^  CaSO  2ao  (solid). 

ZrL2\J  ) 


Plaster  of  paris  is  used  for  making  casts  and  in  surgery.  The 
setting  of  the  material  is  accompanied  by  a  slight  increase  in  volume, 
and  hence  a  very  sharp  reproduction  of  all  the  details  of  the  mold  is 
obtained.  An  "  ivory  "  surface,  which  makes  washing  practicable, 
is  conferred  by  painting  the  cast  with  a  solution  of  paraffin  or  stearin 
in  petroleum  ether.  The  waxy  material,  left  by  evaporation  of  the 
volatile  hydrocarbons,  fills  the  pores  and  prevents  solution  and 
disintegration  of  the  substance  by  water.  Stucco  is  made  with 
plaster  of  paris  and  rubble,  and  is  mixed  with  a  solution  of  glue 
instead  of  water. 

Calcium  Sulphide  CaS.  —  This  compound  is  most  easily  made 
by  strongly  heating  pulverized  calcium  sulphate  and  charcoal.  The 
sulphate  is  reduced: 

4C  +  CaSO4  -»  CaS  +  4CO. 

Calcium  sulphide  is  meagerly  soluble  in  water,  but  is  nevertheless 
slowly  dissolved  in  consequence  of  its  decomposition  by  hydrolysis 
into  calcium  hydroxide  and  calcium  hydrosulphide  Ca(SH)2  (cf.  p. 
254).  Since  calcium  sulphide  is  thus  decomposed  by  water  it  cannot 
be  precipitated  from  aqueous  solution  by  adding  a  soluble  sulphide, 
such  as  ammonium  sulphide,  to  a  solution  of  a  salt  of  calcium.  Only 
the  soluble  hydrosulphide  can  be  formed. 

Ordinary  calcium  sulphide,  after  it  has  been  exposed  to  sunlight, 
usually  shines  in  the  dark.  Barium  sulphide  behaves  in  the  same 


CALCIUM  401 

way.  On  this  account  these  substances  are  used  in  making  luminous 
paint.  They  apparently  owe  this  behavior  to  the  presence  of  traces 
of  compounds  of  vanadium  and  bismuth,  for  the  purified  substances 
are  not  affected  in  the  same  fashion. 

Phosphates  of  Calcium.  —  The  tertiary  orthophosphate  of  cal- 
cium Ca3(PO4)2,  known  as  phosphorite,  is  found  in  many  localities. 
It  is  probably  derived  from  the  remains  of  animals.  Guano  contains 
some  of  the  same  substance,  along  with  compounds  of  nitrogen. 
Apatite,  3Ca3(PO4)2,CaF2,  a  double  salt  with  calcium  fluoride,  is  a 
common  mineral  and  frequent  component  of  rocks.  The  orthophos- 
phate forms  about  83  per  cent  of  bone-ash,  and  is  contained  also  in 
the  ashes  of  plants.  It  may  be  precipitated  by  adding  a  soluble 
phosphate  to  a  solution  of  a  salt  of  calcium. 

Since  it  is  a  salt  of  a  weak  acid,  and  belongs  to  the  less  insoluble 
class  of  such  salts,  calcium  phosphate  is  dissolved  by  dilute  mineral 
acids  (cf.  p.  397),  the  ions  HPO4"  and  H2PO/  being  formed.  When 
a  base,  such  as  ammonium  hydroxide,  is  added  to  the  solution,  the 
calcium  phosphate  is  reprecipitated  (cf.  p.  398). 

Calcium  phosphate  is  chiefly  used  in  the  manufacture  of  phos- 
phorus and  phosphoric  acid  (p.  303),  and  as  a  fertilizer.  The  supply 
of  calcium  phosphate  in  the  soil  arises  from  the  decomposition  of 
rocks  containing  phosphates,  and  is  gradually  exhausted  by  the 
removal  of  crops.  Bone-ash  is  sometimes  used  to  make  up  the 
deficiency.  It  is  almost  insoluble  in  water,  however,  and,  although 
somewhat  less  insoluble  in  natural  water  containing  salts  like  sodium 
chloride,  is  brought  into  a  condition  for  absorption  by  the  plants 
too  slowly  to  be  of  immediate  service.  In  consequence  of  this  the 
"  superphosphate  "  (see  below)  is  preferred. 

Primary  calcium  orthophosphate  ("  superphosphate  ")  is  manu- 
factured in  large  quantities  from  phosphorite  by  the  action  of  sul- 
phuric acid.  The  unconcentrated  "  chamber  acid  "  is  used  for  this 
purpose,  as  water  is  required  in  the  resulting  action.  The  amounts 
of  material  employed  correspond  to  the  equation: 
Ca3(P04)2  +  2H2S04  +  6H20  -» Ca(H2PO4)2,2H2O  +  2CaSO4,2H2O. 

As  soon  as  mixture  has  been  effected,  the  action  proceeds  with  evolu- 
tion of  heat,  and  a  large  cake  of  the  two  hydrated  salts  remains. 
This  mixture,  after  being  broken  up,  dried,  and  packed  in  bags,  is 


402  COLLEGF,    CHEMISTRY 

sold  as  "  superphosphate  of  lime."  The  primary  phosphate  which 
it  contains  is  soluble  in  water,  and  is  therefore  of  great  value  as  a 
fertilizer. 

Calcium  Silicate  CaSiO3.  —  Calcium  metasilicate  CaSi03  forms 
the  mineral  wollastonite,  which  is  rather  scarce,  but  enters  into  the 
composition  of  many  complex  minerals,  such  as  garnet  and  mica. 
It  may  be  made  by  precipitation  with  a  solution  of  sodium  meta- 
silicate (p.  380),  or  by  fusing  together  powdered  quartz  and  calcium 
carbonate: 


Glass.  —  Common  glass  is  a  complex  silicate  of  sodium  and  cal- 
cium, or  a  homogeneous  mixture  of  the  silicates  of  these  metals  with 
silica.  It  has  a  composition  represented  approximately  by  the 
formula  Na2O,CaO,6SiO2;  and  is  made  by  melting  together  sodium 
carbonate,  limestone,  and  pure  sand: 

CaC0   +  6Si0  -»  NaCaCGSiC    +  2C0 


3  2  -  2. 

For  the  most  fusible  glass,  a  smaller  proportion  of  sand  is  employed. 
This  variety  is  known,  from  its  components,  as  soda-glass,  or,  from 
its  easy  fusibility,  as  soft  glass.  Plate-glass  is  made  by  casting  the 
material  in  large  sheets  and  polishing  the  surfaces  until  they  are 
plane.  Window-glass  is  prepared  by  blowing  bulbs  of  long  cylin- 
drical shape,  and  ripping  them  down  one  side  with  the  help  of  a 
diamond.  The  resulting  curved  sheets  are  then  placed  on  a  flat 
surface  in  a  furnace  and  are  there  allowed  to  open  out.  Beads  are 
made,  chiefly  in  Venice,  by  cutting  narrow  tubes  into  very  short 
sections  and  rounding  the  sharp  edges  by  fire.  Ordinary  apparatus 
is  made  of  soft  soda-glass,  and  hence  when  heated  strongly  it  tends 
to  soften  and  also  to  confer  a  strong  yellow  tint  (cf.  p.  380)  on  the 
flame.  Bottles  are  made  with  impure  materials,  and  owe  their 
color  chiefly  to  the  silicate  of  iron  which  they  contain.  In  all  cases 
the  articles  are  annealed  by  being  passed  slowly  through  a  special 
furnace  in  which  their  temperature  is  lowered  very  gradually.  Glass 
which  has  been  suddenly  chilled  is  in  a  state  of  tension  and  breaks 
easily  when  handled. 

Soft  glass  is  partially  dissolved  by  water.    When  powdered  glass 
is  shaken  with  water,  the  liquid  soon  dissolves  enough  sodium 


C  A  LCI  I'M  403 

silicate  to  give  an  alkaline  reaction  with  phenolphthalein  (cf.  p. 
242). 

Bohemian,  or  hard  glass,  is  much  more  difficult  to  fuse  than  soda- 
glass,  and  is  also  much  less  soluble  in  water.  It  is  manufactured  by 
substituting  potassium  carbonate  for  the  sodium  carbonate.  When 
lead  oxide  is  employed  instead  of  limestone,  a  soda-lead  glass  known 
as  flint  glass  is  produced.  This  has  a  high  specific  gravity,  and  a 
great  refracting  power  for  light,  and  is  employed  for  making  glass 
ornaments.  By  the  use  of  grinding  machinery,  cut  glass  is  made 
from  it. 

Colored  glass  is  prepared  by  adding  small  amounts  of  various 
oxides  to  the  usual  materials.  The  oxides  combine  with  the  silica, 
and  produce  strongly  colored  silicates.  Thus,  cobalt  oxide  gives  a 
blue,  chromium  oxide  or  cupric  oxide  a  green,  and  uranium  oxide  a 
yellow  glass.  Cuprous  oxide,  with  a  reducing  agent,  and  compounds 
of  gold,  give  the  free  metals,  suspended  in  colloidal  solution  (p.  96) 
in  the  glass,  and  confer  a  deep-red  color  upon  it.  Milk-glass  contains 
finely  powdered  calcium  phosphate  in  suspension,  and  white  enamels 
are  made  by  adding  stannic  oxide. 

Glass  is  a  typical  amorphous  substance  (pp.  84,  249).  From  the 
facts  that  it  has  no  crystalline  structure,  and  that  it  softens  grad- 
ually when  warmed,  instead  of  showing  a  definite  melting-point,  it 
is  regarded  as  a  supercooled  liquid  of  extreme  viscosity.  Most  single 
silicates  crystallize  easily,  and  have  definite  freezing-  (and  melting-) 
points.  Glass  may  be  regarded  as  a  solution  of  several  silicates. 
When  kept  for  a  considerable  length  of  time  at  a  temperature 
insufficient  to  render  it  perfectly  fluid,  some  of  its  components 
crystallize  out,  the  glass  becomes  opaque,  and  "  devitrification  " 
is  said  to  have  occurred.  The  word  "  crystal  "  popularly  applied 
to  glass  is  thus  definitely  misleading. 

Calcium-ion :  Analytical  Reactions.  —  Ionic  calcium  Ca*  is 
colorless.  It  is  bivalent,  and  combines  with  negative  ions.  Many 
of  the  resulting  salts  are  more  or  less  insoluble  in  water.  Upon  the 
insolubility  of  the  carbonate,  phosphate,  and  oxalate  are  based  tests 
for  calcium-ion  in  qualitative  analysis  (see  p.  440).  The  presence 
of  the  element  is  most  easily  recognized  by  the  brick-red  color  its 
compounds  confer  on  the  Bunsen  flame,  and  by  two  bands  —  a  red 
and  a  green  one  —  which  are  shown  by  the  spectroscope. 


404  COLLEGE    CHEMISTRY 

STRONTIUM  Sr. 

The  compounds  of  strontium  resemble  closely  those  of  calcium, 
both  in  physical  properties  and  in  chemical  behavior. 

Occurrence.  —  The  carbonate  of  strontium  SrC03  is  found  as 
strontianite  (Strontian,  a  village  in  Argyleshire).  The  sulphate, 
celestite  SrSO4,  is  more  plentiful.  The  metal  may  be  isolated  by 
electrolysis  of  the  molten  chloride. 

Compounds  of  Strontium.  —  The  compounds  are  all  made  from 
the  natural  carbonate  or  sulphate.  The  former  may  be  dissolved 
directly  in  acids,  and  the  latter  is  first  reduced  by  means  of  carbon 
to  the  sulphide,  and  then  treated  with  acids. 

Strontium  chloride  SrCl2,6H2O,  made  in  one  of  the  above  ways,  is 
deposited  from  solution  as  the  hexahydrate.  The  nitrate  Sr(N03)2 
comes  out  of  hot  solutions  in  octahedrons  which  are  anhydrous. 
From  cold  water  the  tetrahydrate  Sr(NO3)2,4H2O  is  obtained.  The 
anhydrous  nitrate  is  mixed  with  sulphur,  charcoal,  and  potassium 
chlorate  to  make  "  red  fire."  The  oxide  SrO  may  be  secured  by 
igniting  the  carbonate,  but  it  is  obtained  with  greater  difficulty  than 
is  calcium  oxide  from  calcium  carbonate.  It  is  generally  made  by 
heating  the  nitrate  or  hydroxide. 

Strontium  hydroxide  Sr(OH)2  is  made  by  heating  the  carbonate  in 
a  current  of  superheated  steam : 

SrCO3  +  H2O  ->  Sr(OH)2  +  C02. 

This  action  takes  place  more  easily  than  does  the  mere  dissociation 
of  the  carbonate,  because  the  formation  of  the  hydroxide  liberates 
energy,  and  this  partially  compensates  for  the  energy  which  has  to 
be  provided  to  decompose  the  carbonate.  The  lowering  of  the 
partial  pressure  of  the  carbon  dioxide  by  the  steam  also  contributes 
to  the  result  (cf.  p.  391).  A  hydrate  Sr(OH)2,8H2O  crystallizes  from 
water. 

Strontium-ion  Sr"  is  bivalent,  and  gives  insoluble  compounds  with 
carbonate-ion,  sulphate-ion,  and  oxalate-ion.  The  presence  of 
strontium  is  recognized  by  the  carmine-red  color  which  its  com- 
pounds give  to  the  Bunsen  flame  (see  also  p.  406).  Its  spectrum 
shows  several  red  bands  and  a  very  characteristic  blue  line. 


BARIUM  405 

BARIUM  Ba. 

The  physical  and  chemical  properties  of  the  compounds  of  barium 
recall  those  of  strontium  and  calcium.  All  the  compounds  of  barium 
which  are  soluble  in  water,  or  can  be  brought  into  solution  by  the 
weak  acids  of  the  digestive  fluids,  are  poisonous. 

Occurrence.  —  Like  strontium,  barium  is  found  in  the  form  of 
the  carbonate,  witherite  BaCO3,  and  the  sulphate  BaSO4,  heavy-spar 
or  barite  (Gk.  ftapvs,  heavy).  The  free  metal,  which  is  silver-white, 
may  be  obtained  by  electrolysis  of  the  molten  chloride. 

The  compounds  are  made  by  treating  the  natural  carbonate  with 
acids  directly,  or  by  first  reducing  the  sulphate  with  carbon  to  sul- 
phide, or  converting  the  carbonate  into  oxide,  and  then  treating 
the  products  with  acids. 

Compounds  of  Barium.  —  The  precipitated  form  of  barium 
carbonate  BaC03  is  made  by  adding  sodium  carbonate  to  the  aqueous 
extract  from  crude  barium  sulphide  (q.v.).  Barium  carbonate 
demands  so  high  a  temperature  (about  1500°)  for  the  attainment  of 
a  sufficient  dissociation  tension,  that  special  means  is  employed  for 
its  decomposition.  It  is  heated  with  powdered  charcoal  (cf.  p.  325) : 

BaC03  +  C  -» BaO  +  2CO. 

Natural  barium  sulphate  BaSO4  is  the  source  of  many  of  the  com- 
pounds. The  precipitated  sulphate,  made  by  adding  sulphuric  acid 
to  the  aqueous  extract  from  barium  sulphide,  is  used  in  making 
white  paint  ("  permanent  white  "),in  filling  paper  for  glazed  cards, 
and  sometimes  as  an  adulterant  of  white  lead.  The  salt  is  highly 
insoluble  in  water  and  is  hardly  at  all  affected  by  aqueous  solutions 
of  any  chemical  agents. 

Barium  sulphide  BaS,  like  the  sulphides  of  calcium  and  strontium 
(p.  400),  is  very  slightly  soluble  in  water,  but  slowly  passes  into 
solution  owing  to  hydrolysis  and  formation  of  the  hydroxide  and 
hydrosulphide.  It  is  made  by  heating  the  pulverized  sulphate  with 
charcoal:  BaSO4 -f  4C ->  BaS  +  4CO. 

Barium  chloride  BaCl2,2H2O  is  manufactured  by  heating  the  sul- 
phide with  calcium  chloride.  The  whole  treatment  of  the  heavy- 
spar  is  carried  out  in  one  operation: 

BaSO4  -f  4C  +  CaCl2  -»  4CO  +  BaCl2  -f  CaS. 


406  COLLEGE   CHEMISTRY 

By  rapid  treatment  with  water,  the  chloride  can  be  separated  from 
the  calcium  sulphide  before  much  decomposition  of  the  latter  (cf.  p. 
376)  has  taken  place. 

Barium  chlorate  Ba(C103)2  is  made  by  treating  the  precipitated 
barium  carbonate  with  a  solution  of  chloric  acid.  It  is  deposited  in 
beautiful  monoclinic  crystals,  and  is  used  with  sulphur  and  charcoal 
in  the  preparation  of  "  green  fire." 

Barium  monoxide  BaO  is  manufactured  from  the  carbonate  (see 
above)  or,  in  pure  form,  by  heating  the  nitrate.  The  oxide  unites 
vigorously  with  water  to  form  the  hydroxide. 

The  monoxide,  when  heated  in  a  stream  of  air  or  oxygen,  gives 
barium  peroxide  Ba02,  as  a  compact  gray  mass.  This  change  and 
its  reversal  constitute  the  basis  of  Brin's  process  for  obtaining  oxygen 
from  the  air  (p.  46).  A  hydrate,  BaO2,8H20,  is  thrown  down  as  a 
crystalline  precipitate  when  hydrogen  peroxide  solution  is  added  to 
a  solution  of  barium  hydroxide: 

Ba(OH)2  +  H202  <=±  BaO2  J  +  2H20. 

Barium  peroxide  is  used  in  the  manufacture  of  hydrogen  peroxide. 

Barium  hydroxide  Ba(OH)2,  is  made  by  union  of  the  oxide  with 
water,  or  by  leading  moist  carbon  dioxide  over  the  sulphide  and 
decomposing  the  resulting  carbonate  with  superheated  steam.  It 
is  the  most  soluble  of  the  hydroxides  of  this  group,  and  gives,  there- 
fore, the  highest  concentration  of  hydroxide-ion.  The  solution  is 
known  as  "  baryta-water."  It  is  also  the  most  stable  of  the  three 
hydroxides,  and  may  be  melted  without  decomposition.  A  hydrate 
Ba(OH)2,8H2O  crystallizes  from  water. 

Barium  nitrate  Ba(NO3)2  is  made  by  the  action  of  nitric  acid  on 
the  sulphide,  oxide,  hydroxide,  or  carbonate  of  barium.  The 
crystals  from  aqueous  solution  are  anhydrous. 

Analytical  Reactions  of  the  Calcium  Family.  —  Barium-ion 
Ba*  is  a  colorless,  bivalent  ion.  Many  of  its  compounds  are  insoluble 
in  water,  and  the  sulphate  is  insoluble  in  acids  also.  The  spectrum 
given  by  the  salts  contains  a  number  of  green  and  orange  lines. 

In  solutions  of  salts  of  calcium,  strontium,  and  barium,  the  ions 
may  be  distinguished  by  the  fact  that  calcium  sulphate  solution  will 
precipitate  the  strontium  and  barium  as  sulphates,  but  will  leave 
salts  of  calcium  in  dilute  solution  unaffected.  Similarly,  strontium 


BARIUM  407 

sulphate  solution  precipitates  barium  sulphate,  and  does  not  give 
any  result  with  salts  of  the  first  two.  The  oxalate  of  calcium  is 
precipitated  in  presence  of  acetic  acid,  while  the  oxalates  of  stron- 
tium and  barium  are  not  (cf.  p.  398),  and  there  are  other  differ- 
ences of  a  like  nature  in  the  solubilities  of  the  salts. 

Exercises. —  1.  Arrange  the  chromates  of  the  metals  of  this 
family  in  the  order  of  solubility  (see  Table).  Compare  the  solubili- 
ties with  those  of  the  carbonates,  oxalates,  and  sulphates  of  the 
metals  of  the  same  family. 

2.  What  is  meant  by  fluorescence  (cf.  any  book  on  physics)  ? 

3.  What  will  be  the  ratio  by  volume,  at  150°,  of  the  nitrogen  per- 
oxide and  oxygen  given  off  by  the  decomposition  of  calcium  nitrate? 
What  would  be  the  nature  of  the  difference  between  the  ratio  at  150° 
and  that  at  room  temperature  (cf.  p.  144)? 

4.  Apply  the  rule  of  precipitation  to  the  case  of  adding  sodium 
carbonate  to  a  solution  of  barium  chloride. 

5.  Explain  in  terms  of  the  ionic  hypothesis  the  precipitation  of 
the  sulphate  of  strontium  by  calcium  sulphate  solution,  and  the 
absence  of  precipitation  when  the  latter  is  added  to  a  dilute  solu- 
tion of  a  soluble  salt  of  calcium. 

6.  What  inference  do  you  draw  from  the  fact  that  the  oxalates 
of  barium  and  strontium  are  not  precipitated  in  presence  of  acetic 
acid,  while  the  oxalate  of  calcium  is  so  precipitated?    Is  the  infer- 
ence confirmed  by  reference  to  the  solubility  data? 

7.  Explain  the  fact  that  strontium  and  calcium  chromates  are 
easily  dissolved  by  acetic  acid,  while  barium  chromate  is  dissolved 
only  by  active  mineral  acids. 

8.  Explain  the  fact  that  all  the  carbonates,  save  those  of  potassium, 
sodium,  and  thallium,  are  precipitated  in  neutral  solutions,  but  not 
in  acidified  solutions.     Why  is  the  precipitation  incomplete  when 
carbon  dioxide  is  led  through  solutions  of  salts  of  the  metals,  but 
more  complete  when  the  hydroxides  of  the  metals  are  used? 

9.  Construct  a  table  for  the  purpose  of  comparing  the  properties 
of  the  free  elements  of  this  family  and  also  the  properties  of  their 
corresponding  compounds. 

10.  Are  the  elements  of  this  family  typical  metals  (p.  353)? 


CHAPTER  XXXVI 
COPPER,    SILVER,    GOLD 

THE  three  metals  of  this  family,  being  found  free  in  nature,  are 
amongst  those  which  were  known  in  early  times.  They  are  the 
metals  universally  used  for  coinage  and  for  ornamental  purposes. 
They  are  the  three  best  conductors  of  electricity  (p.  353). 

The  Chemical  Relations  of  the   Copper  Family.  —  Copper 

(Cu,  at.  wt.  63.6),  silver  (Ag,  at.  wt.  107.93),  and  gold  (Au,  at.  wt. 
197.2),  occupy  the  right  side  in  the  second  column  of  the  table  of 
the  periodic  system,  and  the  chemical  relations  of  these  elements 
are  in  many  ways  in  sharp  contrast  to  those  of  the  alkali  metals, 
their  neighbors,  on  the  left  side: 

ALKALI  METALS.  COPPER,  SILVER,  GOLD. 

Very  active;  rapidly  oxidized  by  air;  Amongst  least  active  metals;  only 
displace  all  other  metals  from  com-  copper  is  oxidized  by  air;  displaced 
bination  (E.  M.  series,  p.  245).  by  most  other  metals. 

All  univalent  and  give  but  one  series  Cu1  and  Cn:  two  series.  Ag1:  one 
of  compounds.  Halides  all  soluble  series.  Au1  and  Aum:  two  series, 
in  water.  Chlorides  of  univalent  series  insol. 

Oxides  and  hydroxides  strongly  basic,  Oxides  and  hydroxides  feebly  basic 
and  halides  not  hydrolyzed  (p.353).  (except  Ag2O) ;  halides  hydrolyzed 

(except  Ag-halides).     Hence,  basic 
salts  are  numerous. 

Never  found  in  anion.  Give  no  com-  Frequently  in  anion,  e.g.,  K.Cu(CN)2, 
plex  cations.  K.Ag(CN)2,  K.AuO2,  K.Au(CN),, 

and  in  complex  cations,  e.g., 
Ag(NH3)2.OHandCu(NH3)4.(OH)2. 

On  account  of  their  inactivity  towards  oxygen,  and  their  easy 
recovery  from  combination  by  means  of  heat,  silver  and  gold, 
together  with  the  platinum  family,  are  known  as  the  "  noble  metals." 

COPPER  Cu. 

Chemical  Relations  of  the  Element.  —  Copper  is  the  first 
metallic  element  showing  two  valences  which  we  have  encountered. 
In  such  cases  two  more  or  less  complete,  independent  series  of  salts 
are  known.  These  are  here  distinguished  as  cuprous  (univalent) 

408 


COPPER  409 

and  cupric  (bivalent)  salts.  The  methods  by  which  a  compound  of 
one  series  may  be  converted  into  the  corresponding  compound  of  the 
other  series  should  be  noted. 

The  chief  cuprous  compounds  are  Cu2O,  CuCl,  CuBr,  Cul,  CuCN, 
Cu2S.  There  are  no  cuprous  salts  of  oxygen  acids.  The  cuprous 
compound  is  in  each  case  more  stable  (p.  81)  than  the  corresponding 
cupric  compound,  and  is  formed  from  it  either  by  spontaneous 
decomposition,  as  in  the  cases  of  the  iodide  and  cyanide  (2CuI2 
— •»  2CuI  +  I2),  or  on  heating.  The  cuprous  halides  and  cyanide 
are  colorless  and  insoluble  in  water.  Cuprous-ion  Cu"  seems  to 
be  colorless. 

The  cupric  compounds  are  more  numerous,  as  they  include  also 
salts  of  oxygen  acids,  like  CuSO4,  Cu(NO3)2,  etc.  The  anhydrous 
salts  are  usually  colorless  or  yellow,  but  cupric-ion  Cu"  is  blue,  and 
so,  therefore,  are  the  aqueous  solutions  of  the  salts.  The  cupric  are 
more  familiar  than  the  cuprous  compounds,  since  cupric  oxide, 
sulphate,  and  acetate  are  the  compounds  of  copper  which  most 
frequently  find  employment  in  chemistry  and  in  the  arts.  All  the 
soluble  salts  of  copper  are  poisonous. 

Occurrence.  —  Copper  is  found  free  in  the  Lake  Superior  region, 
in  China,  and  in  Japan.  The  sulphides,  copper  pyrites  CuFeS2  and 
chalcocite  Cu2S,  are  worked  in  Montana,  in  southwest  England,  in 
Spain,  and  in  Germany.  Malachite,  Cu2(OH)2CO3  (=  Cu(OH)2, 
CuC03),  a  basic  carbonate,  is  mined  in  Siberia  and  elsewhere. 
Cuprite  or  ruby  copper  Cu2O  is  also  an  important  ore. 

Extraction  from  Ores.  —  For  isolating  native  copper  it  is  only 
necessary  to  separate  the  metal,  by  grinding  and  washing,  from  the 
rock  through  which  it  ramifies,  and  to  melt  the  almost  pure  powder 
of  copper  with  a  flux  (p.  355).  The  carbonate  and  oxide  ores  require 
coal,  in  addition,  for  the  removal  of  the  oxygen. 

The  liberation  of  copper  from  the  sulphide  ores  is  difficult,  and 
often  involves  very  elaborate  schemes  of  treatment.  This  arises 
from  the  fact  that  many  copper  ores  contain  a  large  amount  of  the 
sulphides  of  iron,  and  these  have  to  be  removed  by  conversion  into 
oxide  (by  roasting)  and  then  into  silicate  (with  sand).  The  silicate 
forms  a  flux,  and  separates  itself  from  the  molten  mixture  of  copper 
and  copper  sulphide.  In  Montana  it  is  found  possible  to  abbreviate 
the  treatment.  The  ore  is  first  roasted  until  partially  oxidized.  It 


410  COLLEGE    CHEMISTRY 

is  then  melted  in  a  cupola  or  a  reverberating  furnace,  and  placed  in 
large  iron  vessels  like  Bessemer  converters  (q.v.)  provided  with  a 
lining  rich  in  silica.  A  blast  of  air  mixed  with  sand  is  next  blown 
through  the  mass.  The  iron  is  completely  oxidized  to  FeO  and 
made  into  silicate  FeSi03,  the  sulphur  escapes  as  sulphur  dioxide, 
and  arsenic  and  lead  are  likewise  removed  by  this  treatment.  The 
silicate  of  iron  floats  as  a  slag  upon  the  copper  when  the  contents  of 
the  converter  are  poured  out.  The  resulting  copper  is  pure  enough 
to  be  cast  in  large  plates  and  purified  by  electrolysis  (see  below). 

Wet  processes  are  used  for  poor  ores.  In  one  of  these  the  ore  is 
roasted  with  salt.  The  copper  is  thus  converted  into  cupric  chloride, 
and  can  be  dissolved  away  from  the  oxide  of  iron  and  other  materials 
by  means  of  water.  Any  traces  of  silver  which  may  have  been 
present  pass  also  into  solution  (as  AgCl),  and  are  precipitated  by 
addition  of  potassium  iodide.  The  copper  is  then  displaced  by 
means  of  scrap  iron,  and  forms  a  brown  sludge  (Cu**  +  Fe  f  —  >  Fe" 


The  tenacity,  ductility,  and  conductivity  of  copper  are  seriously 
affected  by  small  amounts  of  impurities,  such  as  cuprous  oxide  or 
sulphide,  which  are  soluble  in  the  molten  metal.  Hence  a  large 
proportion  of  the  copper  on  the  market  is  purified  by  electrolysis. 
In  this  method  of  refining  the  metal,  thin  sheets  of  copper  form  the 
cathodes,  and  thick  plates  of  copper  the  anodes.  These  are  sus- 
pended alternately  and  close  together  in  large  troughs  filled  with 
cupric  sulphate  solution.  The  cathodes  are  all  connected  with  the 
negative  wire  of  the  dynamo,  and  the  anodes  with  the  positive  one. 
The  Cu"  is  attracted  to  the  cathodes  and  is  deposited  upon  them. 
The  SO4"  migrates  towards  the  anodes,  where  copper  from  the  thick 
plate  becomes  ionized  in  equivalent  amount.  The  stock  of  cupric 
sulphate  thus  remains  the  same,  and  the  practical  effect  of  the 
electrolysis  is  to  carry  copper  across  from  one  plate  to  the  other. 
The  cathodes  are  removed  from  time  to  time,  and  the  deposit  of 
copper  is  stripped  from  their  surface.  Fresh  anodes  are  substituted 
when  the  old  ones  are  eaten  away.  Since  there  is  no  polarization,  a 
current  of  less  than  0.5  volt  suffices. 

Physical  and  Chemical  Properties.  —  Copper  is  red  by  re- 
flected and  greenish  by  transmitted  light.  It  melts  at  1057°,  and 
therefore  much  more  easily  than  pure  iron  (1550°). 


COPPER  411 

In  ordinary  air  copper  becomes  slowly  covered  with  a  green  basic 
carbonate  (not  verdigris,  q.v.).  It  does  not  decompose  water  at  any 
temperature  or  displace  hydrogen  from  dilute  acids  (p.  245).  The 
metal  attacks  oxygen  acids  (pp.  257,  298),  however.  Sea-water 
and  air  slowly  corrode  the  copper  sheathings  of  ships,  giving  a  basic 
chloride,  Cu4(OH)0Cl2,  H20(=  3Cu(OH)2,  CuCl2,  H20);  which  is  found 
in  nature  as  atakamite. 

On  account  of  its  resistance  to  the  action  of  acids,  copper  is  used 
for  many  kinds  of  vessels,  for  covering  roofs  and  ships'  bottoms,  and 
for  coins.  It  furnishes  also  electrotype  reproductions  of  medals,  of 
engraved  plates,  of  type,  etc.  For  this  purpose  a  cast  of  the  object 
is  first  made  in  gutta  percha,  plaster  of  paris,  or  wax.  This  is  then 
coated  with  graphite  to  give  it  a  conducting  surface,  and  receives  an 
electrolytic  deposit  of  copper.  Great  quantities  of  the  metal  are 
used  in  electrical  plants  and  appliances. 

Alloys.  —  The  qualities  of  copper  are  modified  for  special  pur- 
poses by  alloying  it  with  other  metals.  Brass  contains  18-40  per  cent 
of  zinc,  and  melts  at  a  lower  temperature  (p.  390)  than  does  copper. 
A  variety  with  little  zinc  is  beaten  into  thin  sheets,  giving  Dutch- 
metal  ("  gold-leaf  ").  Bronze  contains  3-8  per  cent  of  tin,  11  or 
more  per  cent  of  zinc,  and  some  lead.  It  was  used  for  making  weap- 
ons and  tools  before  means  of  hardening  iron  were  known,  and 
later,  on  account  of  its  fusibility,  continued  to  be  employed  for 
castings  until  displaced  largely  by  cast-iron  (discovered  in  the  eigh- 
teenth century).  Gun-metal  contains  10  per  cent,  and  bell-metal  25 
per  cent  of  tin.  German  silver  contains  19-44  per  cent  of  zinc  and 
6-22  per  cent  of  nickel,  and  shows  none  of  the  color  of  copper. 

Cupric  Chloride  CuClz9  2H2O.  —  This  compound  is  made  by 
union  of  copper  and  chlorine,  or  by  treating  the  hydrate  or  carbonate 
with  hydrochloric  acid.  The  blue  crystals  of  a  hydrate,  CuCl2,2H20, 
are  deposited  by  the  solution.  The  anhydrous  salt  is  yellow.  Dilute 
solutions  are  blue,  the  color  of  cupric-ion,  but  concentrated  solutions 
are  green  on  account  of  the  presence  of  the  yellow  molecules  (p.  227). 
The  aqueous  solution  is  acid  in  reaction  (p.  353).  When  excess  of 
ammonium  hydroxide  is  added  to  the  solution,  the  basic  chloride, 
cupric  oxychloride  Cu4(OH)6Cl2  (above),  which  is  at  first  precipitated, 
redissolves,  and  a  deep-blue  solution  is  obtained  (see  p.  413).  This 


412  COLLEGE   CHEMISTRY 

on  evaporation  yields  deep-blue  crystals  of  hydrated  ammonio-cupric 
chloride  Cu(NH3)4.Cl2,H2O.  The  deep-blue  color  of  the  solution, 
which  is  given  by  all  cupric  salts,  is  that  of  ammonio-cupric-ion 
Cu  (NH3)  4".  The  dry  salt  also  absorbs  ammonia,  giving  CuCl2,6NH3, 
but  a  reduction  of  pressure  results  in  the  final  loss  of  all  the  ammonia. 

Cuprous  Chloride  CuCl.  —  It  may  be  made  by  adding  hydro- 
chloric acid  to  cupric  chloride  solution,  and  boiling  the  mixture  with 
copper  turnings: 

CuCl2  +  Cu  ->  2CuCl    or    Cu"  +  Cu  ->  2Cu\ 

The  solution  contains  compounds  of  cuprous  chloride  with  hydrogen 
chloride  HCl,CuCl  or  HCuCl2  and  H2CuCl3,  which  are  decomposed 
when  water  is  added.  The  cuprous  chloride  is  insoluble  in  water, 
and  forms  a  white  crystalline  precipitate. 

*  The  foregoing  action  is  an  illustration  of  the  fifth  kind  of  ionic 
chemical  change,  namely  that  in  which  a  change  in  valence,  (and  also 
in  the  amount  of  the  electrical  charge),  occurs,  without  any  alteration 
in  the  composition  of  the  ionic  substance.  For  other  illustrations 
see  pp.  Ill  (Mn""  +  4C1'  ->Mn"  +  2C1'  +  C12),  409. 

Cuprous  chloride  is  hydrolyzed  quickly  by  hot  water,  giving, 
finally,  red,  hydrated  cuprous  oxide,  Cu20.  When  dry  it  is  not 
affected  by  light,  but  in  the  moist  state  becomes  violet,  and,  finally, 
nearly  black.  The  action  is  said  to  be  2CuCl  — » CuCl2  +  Cu.  In 
moist  air  it  turns  green,  and  is  oxidized  to  cupric  oxy chloride  (p.  411). 
It  is  dissolved  by  hydrochloric  acid,  giving  the  colorless  complex 
acids  HCuCl2  and  H2CuCl3(see  p.  413).  The  solution  is  oxidized  by 
the  air,  turning  first  brown  and  then  green,  and  finally  depositing 
the  cupric  oxychloride.  Cuprous  chloride  also  dissolves  in  ammo- 
nium hydroxide  (see  p.  413),  giving  ammonio-ciiprous  chloride 
Cu(NH3)2.Cl,  the  ion  Cu(NH3)2'  being  colorless.  The  solution  is 
quickly  oxidized  by  the  air,  turns  deep-blue,  and  then  contains 
Cu(NH3)4". 

The  Bromides  and  Iodides  of  Copper.  —  By  treatment  oi 
copper  with  bromine-water,  and  slow  evaporation  of  the  solution, 
jet-black  crystals  of  anhydrous  cupric  bromide  CuBr2  are  obtained 
(cf.  p.  235).  When  cupric  bromide  is  heated,  bromine  is  given  off, 
and  cuprous  bromide  CuBr  remains. 

Cupric  iodide  CuI2  appears  to  be  unstable  at  ordinary  temperatures, 


COPPER  413 

When  a  soluble  iodide  is  added  to  a  cupric  salt,  a  white  precipitate  of 
cuprous  iodide  Cul  and  free  iodine  are  obtained: 

2Cu"  +  4I7s=>2CuI  J  +  I2. 

The  Solution  of  Insoluble  Salts  when  Complex  Ions  are 
Formed.  —  The  solution  of  an  insoluble  salt  like  cuprous  chloride  by 
hydrochloric  acid  or  ammonium  hydroxide  is  typical  of  a  great 
variety  of  actions  of  which  we  here  meet  with  the  first  examples 
(cf.  p.  253). 

The  dissolving  of  cuprous  chloride  in  hydrochloric  acid  (p.  412), 
to  form  soluble  complex  acids  like  H.CuCL,,  requires  a  special  expla- 
nation. The  complex  negative  ion  CuCl/  which  is  formed,  is  very 
little  dissociated  (CuCL/  <=±  Cu*  +  2C1'),  and  gives  a  smaller  concen- 
tration of  Cu*  than  does  the  insoluble  cuprous  chloride.  Thus,  this 
complex  ion  is  formed  at  the  expense  of  the  Cu"  of  the  insoluble 
cuprous  chloride,  and  the  latter  goes  into  solution  progressively  in 
the  effort  to  restore  the  balance: 

CuCl  (solid)  «=; CuCl  (dslvd)  fc; CV   +  Cu'  )  _^  r  n  ,  ,,  ,    ,. 
2HC1  ^2H*  +  201'!^      Uz   ( 

The  same  exact  laws  of  equilibrium  used  in  discussing  the  dissolving 
of  salts  by  acids  (p.  396)  may  be  applied  to  the  whole  procedure. 

The  dissolving  of  cuprous  chloride  by  the  free  ammonia  of  ammo- 
nium hydroxide  is  explained  in  the  same  way.  The  only  difference 
is  that  here  the  copper  is  in  a  complex  positive  ion.  The  ion 
Cu(NH3)2"  gives  little  Cu"  —  less  than  does  cuprous  chloride,  in  spite 
of  the  insolubility  of  the  latter.  Hence  the  salt  passes  into  solution 
until  the  ion-product  [Cu*]  X  [Cl7],  with  continually  increasing  [Cl7], 
reaches  its  normal  value  or  until  the  solid  is  exhausted. 

The  deep-blue  colored  ion  Cu(NH3)4'*  given  by  cupric  chloride  and 
other  cupric  salts  is  also  very  little  ionized.  Hence  ammonium 
hydroxide  dissolves  all  the  insoluble  cupric  compounds  save  only 
cupric  sulphide,  which  is  the  most  insoluble  of  all  —  that  is,  the  one 
giving  the  smallest  concentration  of  cupric-ion.  Conversely,  the 
sulphide  is  the  only  insoluble  compound  of  copper  which  can  be 
precipitated  from  ammoniacal  solution. 

Cuprous  Oxide  Cu2O.  — This  oxide  is  red  in  color,  and  natural 
specimens  show  octahedral  forms.  It  is  produced  by  oxidation  of 


414  COLLEGE   CHEMISTRY 

finely  divided  copper  at  a  gentle  heat,  or  by  the  addition  of  bases  to 
cuprous  chloride,  and  is  best  made  by  the  action  of  glucose  (p.  332) 
on  cupric  hydroxide.  The  latter  is  reduced  by  the  former,  and 
the  resulting  hydrated  cuprous  oxide  forms  a  yellow  precipitate 
which  quickly  becomes  bright  red.  The  simple  hydroxide,  CuOH, 
is  unknown,  but  the  above  mentioned  precipitate  is  a  hydrated  oxide 
4Cu2O,H2O,  and  yields  Cu2O  when  heated. 

Cuprous  oxide  is  acted  upon  by  hydrochloric  acid,  giving  cuprous 
chloride,  or  rather  HCuCl2.  It  also  dissolves  in  ammonium  hydroxide, 
giving,  probably,  Cu(NH3)2.OH,  which  is  colorless. 

Cupric  Oxide  and  Hydroxide.  —  Cupric  oxide  CuO  is  a  black 
substance  formed  by  heating  copper  in  a  stream  of  oxygen  or  by 
igniting  the  nitrate,  carbonate,  or  hydroxide.  When  heated  strongly 
it  loses  some  oxygen,  and  is  partly  reduced  to  cuprous  oxide. 

Cupric  hydroxide  Cu(OH)2  is  precipitated  as  a  gelatinous  substance 
by  addition  of  sodium  or  potassium  hydroxide  to  a  solution  of  a 
cupric  salt  (Cu"  +  2OH'-»  Cu(OH)2).  When  the  mixture  is  boiled, 
the  hydroxide  loses  water  and  forms  a  black  hydrated  cupric  oxide 
(Cu(OH)2,2CuO  ?).  The  hydroxide  is  soluble  in  ammonium  hydrox- 
ide, with  formation  of  the  compound  Cu(NH3)4.(OH)2,  which  imparts 
a  deep-blue  color  to  the  solution. 


Cupric  Nitrate  Cu(NO3\96H2O.  —  The  nitrate  is  made  by 
treating  cupric  oxide  or  copper  with  nitric  acid  (p.  298),  and  is 
obtained  from  the  solution  as  a  deliquescent,  crystalline  hexahydrate. 
When  dehydrated  at  65°  the  salt  is  partly  hydrolyzed,  and  a  basic 
nitrate  Cu4(OH)6(NO3)2  remains. 

Carbonate  of  Copper.  —  No  normal  carbonate  (CuC03)  can  be 
obtained.  A  basic  carbonate  (malachite)  is  found  in  nature,  and  is 
precipitated  by  adding  soluble  carbonates  to  cupric  salts: 

2CuSO4  +  2Na2CO3  +  H2O  ->  Cu2(OH)2CO3  +  2^80,  +  C02. 
The  carbonate,  if  formed,  would  be  hydrolyzed  by  water  (p.  354). 

Cyanides  of  Copper.  —  With  potassium  cyanide  and  a  solution 
of  a  cupric  salt,  cupric  cyanide  Cu(CN)2  is  precipitated.  This  is  not 
stable,  however,  and  gives  off  cyanogen,  leaving  cuprous  cyanide 
CuCN  :  >  2CuCN  +  C2N2. 


COPPER  415 

Cuprous  cyanide  is  insoluble  in  water,  but  interacts  with  an  excess  of 
potassium  cyanide  solution,  producing  a  colorless  liquid,  from  which 
K.Cu(CN)2  (  =  KCN,CuCN),  potassium  cuprocyanide,  may  be  obtained 
in  colorless  crystals.  The  complex  anion  Cu(CN)/  is  so  little  ionized 
to  Cu"  and  2CN'  that  all  insoluble  copper  compounds,  including 
cupric  sulphide,  are  dissolved  by  potassium  cyanide;  and  none  of 
them  can  be  precipitated  from  the  solution.  Zinc  is  actually  unable 
to  displace  copper  from  such  a  solution.  The  cause  of  the  solution 
of  the  salts  is  the  same  as  when  the  complex  ions  Cu(NH3)2*> 
Cu(NH3)4",  and  CuCl/  are  formed  (p.  413). 

Cupric  Acetate. —  By  the  oxidation  of  plates  of  copper,  sepa- 
rated by  cloths  saturated  with  acetic  acid  (vinegar),  a  basic  acetate 
of  copper  (verdigris)  is  obtained: 

6Cu  +  8HC2H302  +  3O2  ->  2Cu3(OH)2(C2H3O2)4  +  2H20. 

It  is  used  in  manufacturing  green  paint,  is  insoluble  in  water,  and  is 
unaffected  by  light.  It  dissolves  in  acetic  acid,  and  green  crystals 
of  the  normal  acetate  Cu(C2H3O2)2,H2O  are  obtained  from  the  solution. 
The  basic  acetate  is  used  in  preparing  paris  green. 

Cupric  Sulphate  CuSO^  —  This  salt  is  obtained  by  heating 
copper  in  a  furnace  with  sulphur,  and  admitting  air  to  oxidize  the 
cuprous  sulphide.  The  mixture  of  cupric  sulphate  and  cupric  oxide 
which  is  formed  is  treated  with  sulphuric  acid.  The  salt  is  also 
made  by  allowing  dilute  sulphuric  acid  to  trickle  over  granulated 
copper  while  air  has  free  access  to  the  material : 

2Cu  +  2H2S04  +  O2  ->  2CuSO4  +  2H2O. 

This  is  an  example  of  the  use  of  two  reagents  which  separately  have 
little  or  no  action  (cf.  pp.  344-345).  When  concentrated  and  at  a 
high  temperature,  sulphuric  acid  will  itself  act  as  the  oxidizing  agent 
(cf.  p.  263). 

Cupric  sulphate  crystallizes  as  pentahydrate  CuSO4,5H2O  in  blue 
asymmetric  crystals  (Fig.  26,  p.  82),  and  in  this  form  is  called  blue- 
stone  or  blue  vitriol.  The  aqueous  solution  has  an  acid  reaction 
(p.  353).  The  anhydrous  salt  is  white,  and  can  be  crystallized  in 
thin  needles  from  solution  in  hot,  concentrated  sulphuric  acid  (cf. 
p.  82).  Cupric  sulphate  is  employed  in  copper-plating  (p.  410),  in 


416  COLLEGE    CHEMISTRY 

batteries,  as  a  mordant  in  dyeing  (q.v.)  and,  as  a  germicide  and 
insecticide,  for  spraying  plants. 

When  ammonium  hydroxide  is  added  to  cupric  sulphate  solution, 
a  pale-green  basic  sulphate  (Cu4(OH)6SO4?)  is  first  precipitated. 
With  excess  of  the  hydroxide,  the  blue  Cu(NH3)4"  ion  (p.  413)  is 
formed,  and  crystals  of  ammonio-cupric  sulphate  Cu(NH3)4.SO4,H2O 
can  be  obtained  from  the  solution. 

Cupric  sulphate  also  combines  with  potassium  and  ammonium 
sulphates,  giving  double  salts  of  the  form  CuSO4,K2SO4,6H2O,  which 
are  deposited  in  large,  monosymmetric  crystals  from  the  mixed 
solutions.  Double  salts  (p.  231)  exist  as  such  in  the  solid  form  only 
and,  in  water,  are  resolved  into  their  components. 

The  Sulphides  of  Copper.  —  Cuprous  sulphide  Cu2S  occurs  in 
nature  in  rhombic  crystals  of  a  gray,  metallic  appearance.  It  is 
made  by  heating  cupric  sulphide,  a  stream  of  hydrogen  gas  being  used 
to  assist  the  removal  of  the  excess  of  sulphur. 

Cupric  sulphide  CuS  is  deposited  as  a  black  precipitate  when 
hydrogen  sulphide  is  led  through  a  solution  of  a  cupric  salt.  Made 
in  this  way,  it  is  always  partly  decomposed  into  Cu2S  +  S.  By 
cautiously  treating  copper  with  excess  of  sulphur  at  114°  it  may  be 
obtained  as  a  blue  crystalline  solid.  At  higher  temperatures  it 
gives  off  sulphur. 

Analytical  Reactions  of  Compounds  of  Copper.  —  The  ion 

of  ordinary  cupric  salts,  cupric-ion  Cu",  is  blue,  and  that  of  cuprous 
salts,  cuprous-ion  Cu*,  is  colorless.  Cuprous  solutions,  however,  are 
easily  oxidized  by  the  air  and  become  blue.  In  solutions  containing 
cupric-ion,  hydrogen  sulphide  precipitates  cupric  sulphide,  even  in 
presence  of  acids  (p.  398).  Bases  throw  down  the  blue  hydroxide, 
and  carbonates  precipitate  a  green  basic  salt  (p.  414).  Potassium 
ferrocyanide  gives  the  brown,  gelatinous  cupric  ferrocyanide, 

2Cu.SO4  +  K4.Fe(CN)6^±Cu2.Fe(CN)6l  +  2K2S04. 

A  characteristic  test  is  the  formation  of  the  deep-blue  Cu(NHa)4'* 
ion  with  excess  of  ammonium  hydroxide.  This  solution  itself  re- 
sponds to  certain  precipitants  (e.g.  H2S)  only.  Solutions  of  complex 
cuprous  and  cupric  cyanides  such  as  K.Cu(CN)2  and  K2Cu(CN)4 
are  colorless,  and  do  not  respond  to  any  of  the  above  tests.  With 


SILVER  417 

mierocosmic  salt  or  borax  (pp.  312,  350),  copper  compounds  form  a 
bead  which  is  blue  in  the  oxidizing  part  of  the  flame  and  becomes 
red  and  opaque  (liberation  of  copper)  in  the  reducing  flame. 

SILVER. 

Chemical  Relations  of  the  Element.  —  This  element  presents 
a  curious  assortment  of  chemical  properties.  It  differs  from  copper 
in  having  a  strongly  basic  oxide,  and  in  giving  salts  with  active  acids 
which  are  not  hydrolyzed  by  water.  In  these  respects  it  approaches 
the  metals  of  the  alkalies  and  alkaline  earths.  It  resembles  copper 
in  entering  into  complex  compounds,  and  in  giving  insoluble  halides. 
It  differs  from  both  copper  and  the  metals  of  the  alkalies,  and 
resembles  gold  and  platinum,  in  that  its  oxide  is  easily  decomposed 
by  heat,  with  formation  of  the  free  metal,  and  in  the  low  position  it 
occupies  in  the  electromotive  series  and  the  consequent  slight  chem- 
ical activity  of  the  free  metal. 

Occurrence.  —  Native  silver,  usually  scattered  through  a  rocky 
matrix,  contains  varying  amounts  of  gold  and  copper.  Native 
copper  always  contains  dissolved  silver.  Sulphide  of  silver  (Ag2S) 
occurs  alone  and  dissolved  in  galenite  (PbS).  Smaller  amounts  of 
the  metal  are  obtained  from  pyrargyrite  Ag3SbS3  and  proustite 
Ag3AsS3,  and  from  horn-silver  AgCl. 

Metallurgy.  — The  silver  contained  free,  or  as  sulphide,  in  ores  of 
copper  and  lead,  is  found  in  the  free  state  dissolved  in  the  metals 
extracted  from  these  ores,  and  is  secured  by  refining  them.  In  the 
electrolytic  refining  of  copper,  silver  is  obtained  from  the  mud 
deposited  in  the  baths  (p.  410).  The  proportion  present  in  lead  is 
usually  small.  Parke's  process,  by  which  the  silver  is  separated  from 
the  lead,  takes  advantage  of  the  fact  that  molten  zinc  and  lead  are 
practically  insoluble  in  one  another,  while  silver  is  much  more 
soluble  in  zinc  than  in  lead.  Lead  dissolves  1.6  per  cent  of  zinc,  and 
zinc  1.2  per  cent  of  lead.  The  principle  is  the  same  as  in  the  removal 
of  iodine  from  water  by  ether  (p.  103).  The  lead  is  melted  and 
thoroughly  mixed  by  machinery  with  a  small  proportion  of  zinc. 
After  a  short  time  the  zinc  floats  to  the  top,  carrying  with  it  in 
solution  almost  all  of  the  silver,  and  solidifies  at  a  temperature  at 
which  the  lead  is  still  molten.  The  zinc-silver  alloy  is  skimmed  off, 


418  COLLEGE    CHEMISTRY 

and  heated  moderately  in  a  furnace  to  permit  the  adhering  lead  to 
drain  away.  The  zinc  is  finally  distilled  off  in  clay  retorts,  and  the 
lead  remaining  with  the  silver  is  removed  by  cupellation.  This 
operation  consists  in  heating  the  molten  metal  strongly  in  a  blast  of 
air.  The  lead  is  converted  into  litharge  (PbO) ,  which  flows  in  molten 
condition  over  the  edge  of  the  cupel. 

Ores  of  silver  which  do  not  contain  much  or  any  lead  are  often 
smelted  with  lead  ores,  and  the  product  is  treated  as  described 
above,  but  many  other  processes  are  in  use.  Thus,  sulphide  ores 
are  sometimes  roasted  until  the  iron  and  part  of  the  copper  are 
converted  into  oxide  while  the  rest  of  the  copper  ,and  all  the 
silver  remain  as  sulphate.  The  metal  is  secured  by  extracting  the 
mass  with  water  and  precipitating  the  silver  by  means  of  copper 
(p.  245).  Some  ores  are  roasted  with  salt,  and  the  resulting 
chloride  of  silver  is  dissolved  out  with  sodium  thiosulphate,  or  even 
strong  brine. 

During  the  first  half  of  the  nineteenth  century  the  world's  total 
output  of  silver  averaged  only  643  tons  per  year.  Up  to  1870  a 
gram  of  gold  could  buy  15.5  g.  of  silver.  Now  that  the  produc- 
tion has  reached  6000  tons,  the  same  amount  of  gold  purchases 
about  35  g. 

Physical  Properties.  —  Pure  silver  is  almost  perfectly  white. 
It  melts  at  960°.  Its  ductility  is  so  great  that  wires  can  be  drawn  of 
such  fineness  that  2  kilometers  weigh  only  about  1  g.  In  the  molten 
condition  it  absorbs  mechanically  about  twenty-two  times  its  own 
volume  of  oxygen,  but  gives  up  almost  all  of  this  as  it  solidifies. 
Fantastically  irregular  masses  result  from  the  "  sprouting "  or 
"  spitting  "  which  accompanies  the  escape  of  the  gas. 

By  addition  of  ferrous  citrate  to  silver  nitrate,  a  red  solution  and 
lilac  precipitate  of  free  silver  can  be  made.  The  latter,  after  washing 
with  ammonium  nitrate  solution,  gives  a  red,  colloidal  solution  in 
water.  Such  colloidal  solutions  of  metals  are  formed  also  by  passing 
an  electrical  discharge  between  wires  of  silver,  gold,  or  platinum 
held  under  water. 

Silver  is  alloyed  with  copper  to  render  it  harder.  The  silver  coin- 
age of  the  United  States  and  the  continent  of  Europe  has  a  "  fine- 
ness of  900  "  (900  parts  of  silver  in  1000),  and  that  of  Great  Britain 
925.  Silver  ornaments  have  a  fineness  of  800  or  more. 


SILVER  419 

Chemical  Properties.  —  Silver  does  not  combine,  with  oxygen 
either  in  the  cold  or  when  heated.  It  does  not  ordinarily  displace 
hydrogen  from  aqueous  solutions  of  acids.  Silver  interacts  with 
cold  nitric  acid  and  with  hot,  concentrated  sulphuric  acid,  giving 
the  nitrate  or  sulphate  of  silver  and  oxides  of  nitrogen  or  of  sulphur 
(pp.  257,  298). 

The  Halides  of  Silver.  —  The  chloride  AgCl,  bromide  AgBr,  and 
iodide  Agl  are  formed  as  curdy  precipitates  when  a  salt  of  silver  is 
added  to  a  solution  containing  the  appropriate  halide  ion.  The  first 
is  white,  and  melts  at  about  457°.  The  second  and  third  are  very 
pale-yellow  and  yellow  respectively.  The  insolubility  in  water, 
which  is  very  great,  increases  in  the  above  order. 

When  exposed  to  light,  the  chloride  becomes  first  violet  and  finally 
brown,  chlorine  being  liberated.  The  bromide  and  iodide  behave 
similarly.  Solid  silver  chloride  absorbs  ammonia,  forming  first 
2AgCl;3NH3,  and  then  AgCl,3NH3. 

Complex  Compounds  of  Silver.  —  Silver  chloride  dissolves 
easily  in  excess  of  ammonium  hydroxide,  giving  the  complex  cation 
Ag(NH3)2".  The  bromide,  which  is  less  readily  soluble,  gives  the 
same  complex  ion.  The  iodide  is  hardly  soluble  at  all.  Ammonio- 
argentic-ion  Ag(NH3)2",  in  solutions  of  concentrations  such  as  are 
commonly  used  (.1  N  to  N),  gives  about  the  same  concentration  of 
argention  Ag*  as  does  the  bromide,  and  much  more  than  the  highly 
insoluble  iodide  (cf.  p.  413).  Hence  the  latter  is  almost  insoluble  in 
ammonium  hydroxide,  and  can  be  precipitated  in  ammoniacal  solu- 
tion. All  three  of  the  insoluble  halides  dissolve  in  solutions  of  potas- 
sium cyanide  and  of  sodium  thiosulphate,  as  do  also  all  the  other 
insoluble  silver  salts.  Usually  an  equivalent  amount  of  the  cyanide 
or  thiosulphate  suffices,  but  for  solution  of  the  sulphide  an  excess 
is  required.  With  the  cyanide,  double  decomposition  gives  first  the 
insoluble  silver  cyanide  (AgCN)  which  then  dissolves,  forming  the 
soluble  potassium  argenti-cyanide  K.  Ag(CN)2.  The  thiosulphate  gives 
a  solution  containing  the  complex  salt  Na3.Ag(S2O3)2.  The  more 
active  metals,  like  zinc  and  copper,  displace  silver  from  all  solutions, 
whether  the  solutions  contain  simple  or  complex  salts. 

Oxides  of  Silver.  —  When  sodium  hydroxide  is  added  to  a 
solution  of  a  salt  of  silver,  a  pale-brown  precipitate  is  obtained, 


420  COLLEGE    CHEMISTRY 

which,  after  being  freed  from  water,  is  found  to  be  argentic  oxide 
Ag2O,  and  not  AgOH.  The  aqueous  solution  of  argentic  oxide,  how- 
ever, is  distinctly  alkaline,  and  presumably  therefore  does  contain 
the  hydroxide:  2AgOH  <±  Ag2O  +  H20. 

Argentic  oxide  melts  and  gives  off  its  oxygen  at  250-270°.  It  is  an 
active  basic  oxide.  When  moist,  it  absorbs  carbon  dioxide  from  the 
air.  With  ammonium  hydroxide  it  forms  the  soluble  Ag(NH4)2.OH. 

Silver  peroxide  Ag2O2  is  formed  by  the  action  of  ozone  on  silver. 
In  the  electrolysis  of  silver  nitrate  it  is  deposited  in  shining  black 
crystals  on  the  anode.  There  is  also  a  suboxide  Ag4O. 

Salts  of  Silver.  —  Silver  nitrate  AgNO3  is  obtained  by  treating 
silver  with  aqueous  nitric  acid: 

3Ag  +  4HNO3  -»  3AgNO3  +  NO  +  2H2O. 

From  the  solution,  colorless  rhombic  crystals  (Fig.  7,  p.  8)  are 
deposited.  These  melt  at  218°.  In  the  form  of  thin  sticks  made  by 
casting  (lunar  caustic),  the  substance  is  used  in  medicine,  partly 
because  it  combines  with  albumins  to  form  insoluble  compounds. 
The  aqueous  solution  is  neutral.  The  pure  salt  is  not  affected  by 
light,  but  when  deposited  on  cloth,  on  the  skin  of  the  fingers,  or  on 
the  mouth  of  the  reagent  bottle,  it  is  converted  into  the  chloride, 
and  from  this,  in  turn,  silver  is  liberated.  For  this  reason  it  is  an 
ingredient  in  some  marking-inks. 

Silver  carbonate,  the  neutral  salt  Ag2CO3,  and  not  a  basic  carbonate, 
is  precipitated  from  solutions  of  salts  of  silver  by  soluble  carbonates. 
It  is  slightly  yellow  in  color.  With  water  it  gives  a  faint  alkaline 
reaction,  and,  like  calcium  carbonate,  is  soluble  in  excess  of  carbonic 
acid  (p.  391).  When  heated,  the  carbonate  decomposes,  leaving 
metallic  silver.  The  sulphate  Ag3S04  is  made  by  the  action  of  con- 
centrated sulphuric  acid  on  the  metal.  When  it  is  mixed  with  a 
solution  of  aluminium  sulphate  (q.v.),  octahedral  crystals  of  silver- 
alum  Ag2SO4,Al2(SO4)3,24H2O  are  obtained.  Silver  sulphide  Ag2S  is 
precipitated  by  hydrogen  sulphide  from  solutions  of  all  silver  com- 
pounds, whether  free  acids  are  present  or  not,  and  irrespective  of 
the  form  in  which  the  silver  is  combined.  Excess  of  potassium 
cyanide,  however,  prevents  its  precipitation  from  the  argenticyanide. 
The  sulphide  is  formed  by  the  action  of  metallic  silver  on  alkaline 
hydrosulphides,  and  this  interaction  forms  the  basis  of  the  "  hepar  " 


SILVER  421 

test  for  sulphur.  Silver  orthophosphate  Ag3PO4  (yellow),  arsenate 
Ag3AsO4  (brown),  and  chromate  Ag2CrO4  (crimson),  are  produced 
by  precipitation,  and  their  distinctive  colors  enable  us  to  use  silver 
nitrate  in  analysis  as  a  reagent  for  identifying  the  acid  radicals. 

Electroplating.  —  The  process  is  similar  to  the  electro-deposition 
of  copper  (p.  410).  The  article  to  be  plated  is  cleaned  with  extreme 
care  and  attached  to  the  negative  wire.  A  plate  of  silver  forms  the 
positive  electrode  and,  since  simple  salts  of  silver  do  not  give  coherent 
deposits,  the  bath  is  a  solution  of  potassium  argenticyanide.  The 
potassium-ion  (K*)  migrates  to  the  negative  wire  and,  since  potassium 
requires  a  much  greater  E.M.F.  for  its  liberation  than  does  silver, 
silver  is  there  deposited  from  the  trace  of  argentic-ion  given  by  the 
complex  silver  ions  in  the  neighborhood: 

Ag(CN)/  1=5  Ag-  +  2CN',      Ag'  +  Q  -»  Ag. 

At  the  positive  electrode  silver  goes  into  solution  in  equivalent 
amount,  giving  argentic-ion,  and  the  above  equations  are  reversed. 
Mirrors  are  silvered  through  the  reduction  of  silver  nitrate  by 
organic  compounds  such  as  potassium-sodium  tartrate  (Rochelle 
salt),  glycerine,  formaldehyde  (formalin),  or  grape  sugar. 

Photography.  —  Bromo-gelatine  dry  plates  are  covered  with 
an  emulsion  of  gelatine  in  which  silver  bromide  is  suspended. 

After  exposure,  often  for  only  a  fraction  of  a  second,  there  is  no 
visible  alteration  in  the  film.  The  image  is  developed.  Chemically, 
this  consists  in  reducing  the  silver  bromide  to  metallic  silver  by 
means  of  reducing  agents.  While  the  whole  of  the  halide  upon  the 
plate  is  reducible,  if  the  reducing  agent  is  kept  upon  it  for  a  sufficient 
length  of  time,  the  parts  reached  by  the  light  are  affected  first,  and 
with  a  speed  proportional  to  the  intensity  of  the  illumination  under- 
gone by  each  part.  The  unreduced  silver  bromide  is  then  dissolved 
out  with  sodium  thiosulphate  ("  hyposulphite  of  soda  "  or  "  hypo  "), 
and  the  silver  image  is  thus  saved  from  being  fogged  over  by  the 
silver  that  would  be  deposited  if  the  plate  were  to  be  brought  into  the 
light  without  this  treatment  (fixing).  The  result  is  a  "  negative,"  as 
the  parts  brightest  in  the  object  are  now  opaque,  and  the  darkest 
parts  of  the  object  are  transparent. 

The  simplest  developer  is  potassium-ferrous  oxalate  K2.Fe(C2O4)2, 
a  solution  of  which  may  be  made  by  mixing  ferrous  sulphate  and 


422  COLLEGE   CHEMISTRY 

potassium  oxalate.  For  the  sake  of  simplicity  we  may  regard  the 
action  as  a  reduction  by  means  of  ferrous  oxalate,  which  itself  is 
oxidized  to  ferric  oxalate  (Fe2(C2O4)3): 

3FeC2O4  +  3Ag4Br-»  Fe^OJ,  +  FeBrs  +  3Ag. 

In  brief,  we  have  3Fen  becoming  2Fem  +  Fenl,  and  this  amount  of 
bivalent  iron  therefore  takes  up  3Br,  liberating  the  silver  with  which 
it  was  combined.  The  developers  commonly  employed  are  alkaline 
solutions  of  the  sodium  salts  of  hydroquinone  and  pyrogallic  acid. 

In  printing,  the  light  and  dark  are  again  reversed,  the  denser  parts 
of  the  negative  protecting  the  compounds  on  the  paper  below  it  from 
action,  and  leaving  them  white.  Either  "  bromide  "  papers,  which 
require  only  brief  exposure  and  are  developed  like  the  plate,  are  used, 
or  silver  chloride  is  the  sensitive  substance,  and  prolonged  exposure 
to  light  is  allowed  to  liberate  the  proper  amount  of  silver.  The 
operation  of  fixing  is  performed  as  before.  In  toning,  a  solution  of 
sodium  chloraurate  is  employed.  A  portion  of  the  silver  dissolves, 
displacing  gold  (p.  245),  which  is  deposited  in  its  place: 

NaAuCl4  +  3Ag  ->  NaCl  +  3AgCl  +  Au. 
The  thin  film  of  gold  gives  a  richer  color  to  the  print. 

Analytical  Reactions  of  Silver  Compounds. — Argentic-ion 
Ag*  is  colorless.  Many  of  its  compounds  are  insoluble,  the  precipi- 
tation of  the  chloride,  which  is  insoluble  in  dilute  acids,  being  used 
as  a  test.  Mercurous  chloride  and  lead  chloride  are  also  white  and 
insoluble,  but  silver  chloride  dissolves  in  ammonium  hydroxide, 
mercurous  chloride  (q.v.)  turns  black,  and  lead  chloride  is  not  altered 
in  color  (and  is  also  soluble  in  hot  water).  With  excess  of  ammonium 
hydroxide,  silver  salts  give  the  complex  cation  Ag(NH3)2*  and,  from 
solutions  containing  silver  in  this  form,  only  the  iodide  and  sulphide 
can  be  precipitated.  Sodium  thiosulphate  and  potassium  cyanide 
dissolve  all  silver  salts,  giving  salts  of  complex  acids  with  silver  in 
the  anion  (p.  419). 

GOLD. 

Chemical  Relations  of  the  Element.  —  This  element  forms 
two  very  incomplete  series  of  compounds  corresponding  respectively 
to  aurous  and  auric  oxides,  Au2O  and  Au203.  The  former  is  a 
feebly  basic  oxide,  the  latter  mainly  acid-forming.  No  simple  salts 


GOLD  423 

with  oxygen  acids  are  stable.  All  the  compounds  of  gold  are  easily 
decomposed  by  heat  with  liberation  of  the  metal.  All  other  common 
metals  displace  gold  from  solutions  of  its  compounds  (p.  245).  Mild 
reducing  agents  likewise  liberate  gold.  The  element  enters  into 
many  complex  anions. 

Occurrence  and  Metallurgy.  —  Gold  is  found  chiefly  in  the  free 
condition  disseminated  in  veins  of  quartz,  or  mixed  with  alluvial 
sand.  Small  quantities  are  found  also  in  sulphide  ores  of  iron  and 
copper.  Telluride  of  gold  (sylvanite),  in  which  silver  takes  the 
place  of  a  part  of  the  gold  [Au,Ag]Te2,  is  found  in  Colorado. 

From  the  alluvial  deposits,  gold  is  usually  separated  by  washing 
in  a  cradle,  as  in  the  Klondyke.  Quartz  veins,  which  in  the  Trans- 
vaal Colony  reach  a  thickness  of  a  meter  and  carry  an  average  of 
18  g.  of  gold  per  ton,  are  mined,  and  the  material  is  pulverized  with 
stamping  machinery.  About  55  per  cent  of  the  gold  is  then  sepa- 
rated by  allowing  the  powdered  rock  to  be  carried  by  a  stream  of 
water  over  copper  plates  amalgamated  with  mercury.  The  gold 
dissolves  in  the  latter,  and  is  secured  by  removal  and  distillation  of 
the  amalgam.  The  finer  particles,  contained  in  the  sludge  which 
runs  off  ("  tailings  "),  are  extracted  by  adding  a  dilute  solution  of 
potassium  cyanide  (MacArthur-Forest  process)  and  exposing  the 
mixture  to  the  air.  Oxidation  and  simultaneous  interaction  with 
the  cyanide  give  potassium  aurocyanide  KAu(CN)2.  From  this 
solution  the  gold  is  isolated,  either  by  electrolysis,  or  in  the  form  of 
a  purple  powder  by  precipitation  with  zinc. 

Auriferous  pyrites  is  roasted,  and  then  treated  with  chlorine  gas. 
The  chloride  of  gold  which  is  formed  is  dissolved  out  with  water. 
From  the  solution,  the  gold  is  precipitated  with  ferrous  sulphate  or 
oxalic  acid: 

2AuClt  +  6FeSOi->2Fea(S04)3  +  2FeCl,  +  2Au. 

In  the  former  case  a  purple  powder,  and  in  the  latter,  if  the  solution 
is  heated,  a  spongy  mass  (the  form  used  by  dentists),  is  obtained. 

The  gold  separated  from  ores  in  the  above  ways  contains  silver, 
copper,  lead,  and  other  metals,  and  various  methods  of  refining, 
electrolytic  and  otherwise,  are  used.  In  one  of  these  the  gold  is 
melted,  and  a  stream  of  chlorine  is  passed  through  it.  The  metals, 
excepting  gold,  are  converted  into  chlorides.  The  chloride  of  silver 


424  COLLEGE   CHEMISTRY 

rises  as  a  liquid  to  the  surface,  while  chlorides  of  arsenic  and  anti- 
mony are  volatilized.  A  layer  of  melted  borax  prevents  loss  of 
silver  chloride  by  volatilization.  The  silver  chloride,  when  it  has 
solidified,  is  placed  between  wrought-iron  plates  and  immersed  in  an 
electrolyte  (usually  dilute  sulphuric  acid). 

The  world's  production  of  gold  during  the  first  half  of  the  nine- 
teenth century  averaged  27  tons  annually.  In  1897  it  was  363  tons, 
and  in  1899,  472.6  tons.  In  the  former  year  North  America,  includ- 
ing Canada,  produced  28.5  per  cent  of  the  whole,  the  Transvaal 
Colony  23.2  per  cent,  and  Australia  21.2  per  cent. 

Properties  of  the  Metal.  —  Gold  is  yellow  in  color,  and  is  the 
most  malleable  and  ductile  of  all  the  metals.  It  melts  at  1064°.  To 
give  it  greater  hardness  it  is  alloyed  with  copper,  the  proportion  of 
gold  being  defined  in  "  carats."  Pure  gold  is  "  24-carat."  British 
sovereigns  are  22-carat  and  contain  ^  of  copper.  American,  French, 
and  German  coins  are  21.6-carat,  or  90  per  cent  gold. 

Gold  is  not  affected  by  free  oxygen  nor  by  hydrogen  sulphide.  It 
does  not  displace  hydrogen  from  dilute  acids,  nor  does  it  interact 
with  nitric  or  sulphuric  acids  or  any  oxygen  acids  except  selenic  acid. 
It  combines,  however,  with  free  chlorine,  and  it  therefore  interacts 
with  a  mixture  of  nitric  and  hydrochloric  acids  (aqua  regia),  which 
gives  off  this  gas  (p.  299).  Chlorauric  acid  H.AuCl4(=  HCl,AuCl3) 
is  formed,  and  the  action  is  assisted  by  the  fact  that  the  gold  ions 
are  taken  into  the  little-dissociated  anion  AuCl/.  Gold  is  the  least 
active  of  the  familiar  metals. 

Compounds  with  the  Halogens.  —  Chlorauric  acid,  formed  as 
above,  is  deposited  in  yellow,  deliquescent  crystals  of  H.  AuCl4,4H2O. 
The  yellow  sodium  chloraurate  NaAuCl4,2H2O,  obtained  by  neutrali- 
zation of  the  acid,  is  used  in  photography  (p.  422).  The  acid  gives 
up  hydrogen  chloride  when  heated  very  gently,  leaving  the  red, 
crystalline  auric  chloride  AuCl3.  When  auric  chloride  is  heated  to 
180°  aurous  chloride  AuCl  and  chlorine  are  formed.  This  salt  is  a 
white  powder.  It  is  insoluble  in  water,  but  in  boiling  water  is  con- 
verted quickly  into  auric  chloride  and  free  gold:  3 AuCl— >AuCl34-2Au. 

Other  Compounds.  —  When  caustic  alkalies  are  added  to  chlor- 
auric  acid,  or  to  sodium  chloraurate,  auric  hydroxide  Au(OH)3  is  pre- 
cipitated. This  substance  is  an  acid,  and  interacts  with  excess  of 


GOLD  425 

the  base,  forming  aurates.  These  are  derived  from  met-auric  acid 
(Au(OH)3  —  H2O  =  HAuO2),  as,  for  example,  potassium  aurate 
K.AuO2,3H2O.  This  salt  interacts  by  double  decomposition,  giving, 
for  instance,  with  silver  nitrate,  the  insoluble  silver  salt  AgAuO2. 
Its  solution  is  alkaline  in  reaction,  showing  that  auric  acid  is  a  weak 
acid  (cf.  p.  353). 

Auric  oxide  Au2O3  is  a  brown,  and  aurous  oxide  Au2O  is  a  violet 
powder. 

On  account  of  its  reducing  action,  hydrogen  sulphide  precipitates 
from  chlorauric  acid  a  dark-brown  mixture  containing  much  aurous 
sulphide  Au2S  and  free  sulphur,  as  well  as  some  auric  sulphide  Au2S3. 

The  aurocyanides,  like  K.Au(CN)2(=  KCN,AuCN),  and  the  auri- 
cyamdes,  like  K.Au(CN)4(==KCN,Au(CN)3),  are  formed  by  the  action 
of  potassium  cyanide  on  aurous  and  auric  compounds,  respectively. 
They  are  colorless  and  soluble.  Their  solutions  are  used  as  baths, 
in  conjunction  with  a  gold  anode,  for  electrogilding. 

Assaying. — In  assaying,  the  material  containing  the  gold  is  heated 
with  borax  and  lead  in  a  small  crucible  (cupel)  of  bone-ash.  The 
lead  and  copper  are  oxidized,  and  the  oxides  are  absorbed  by  the 
cupel,  leaving  a  drop  of  molten  alloy  of  gold  and  silver.  The  cold 
button  is  flattened  by  hammering  and  rolling,  and  treated  with 
nitric  acid  to  remove  the  silver.  The  gold,  which  remains  un- 
attacked,  is  washed,  fused  again,  and  weighed.  The  acid  will  not 
interact  with  the  silver  and  remove  it  completely  if  the  quantity  of 
gold  exceeds  25  per  cent.  When  the  proportion  of  gold  is  greater 
than  this,  a  suitable  amount  of  pure  silver  is  fused  with  the  alloy 
("  quartation  "). 

Exercises.  —  1.  Write  equations  for  the  interactions,  (a)  of  salt 
water  and  oxygen  with  copper  (p.  411),  (6)  of  ferrous  oxide  and  sand 
(p.  410). 

2.  Write  the  formulae  of  the  basic  chloride,  nitrate,  carbonates, 
and  sulphate  of  copper  as  if  these  substances  were  composed  of  the 
normal  salt,  the  oxide  and  water  (p.  411). 

3.  Can  you  develop  any  relation  between  the  facts  that  solutions 
of  cupric  salts  are  acid  in  reaction  and  that  they  give  basic  carbonates 
by  precipitation? 

4.  Formulate  the  action  of  potassium  cyanide  in  dissolving  cupric 


426  COLLEGE   CHEMISTRY 

hydroxide  and  cuprous  sulphide,  assuming  that  potassium  cupro- 
cyanide  is  formed. 

5.  How  should  you  set  about  making  cupric  orthophosphate  (in 
solution),  ammonium  cuprocyanide,  and  lead  cuprocyanide? 

6.  Write  the  formulae  of  some  of  the  double  salts  analogous  to 
potassium-cupric  sulphate  (p.  416). 

7.  What  chemical   agents  are  present  in  a  Bunsen  flame?    If 
borax  beads  were  made  in  the  oxidizing  flame  with  cupric  chloride, 
cuprous  bromide,  and  cupric  sulphate,  severally,  what  actions  would 
take  place? 

8.  Write  the  equations  for  the  interaction  of,  (a)  silver  and  con- 
centrated sulphuric  acid,  (b)  silver  chloride  and  sodium  carbonate 
when  heated  strongly,  (c)  sodium  thiosulphate  and  silver  bromide. 

9.  What  reagents  should  you  use  to  precipitate  the  phosphate, 
arsenate,  and  chromate  of  silver? 

10.  Write  the  equations  for  the  interactions  of,  (a)  potassium 
hydroxide  and  auric  hydroxide,  (6)  potassium  cyanide  and  sodium 
chloraurate. 

11.  In  what  respects  are  the  elements  of  this  family  distinctly 
metallic,  and  in  what  respects  are  they  allied  to  the  non-metals 
(p.  353)? 

12.  Collect  all  the  evidence  tending  to  show  that  the  cuprous 
compounds  are  more  stable  than  the  cupric. 

13.  Make  a  classified  list  of  the  methods  by  which  cupric  com- 
pounds are  transformed  into  cuprous,  and  vice  versa. 

14.  Of  which  metals  should  it  be  possible  to  obtain  colloidal 
solutions,  and  of  which  not  (p.  245)  ? 


CHAPTER  XXXVII 

GLUCINUM,    MAGNESIUM,    ZINC,    CADMIUM,    MERCTTBY. 

THE    RECOGNITION    OF    CATIONS    IN   QUALITATIVE 

ANALYSIS 

The  Chemical  Relations  of  the  Family.  —  The  remaining 
elements  of  the  third  column  of  the  periodic  table,  namely,  glucinum 
or  beryllium  (Gl,  or  Be,  at.  wt.  9.1),  magnesium  (Mg,  at.  wt.  24.36), 
zinc  (Zn,  at.  wt.  65.4),  cadmium  (Cd,  at.  wt.  112.4),  and  mercury  (Hg, 
at.  wt.  200.0),  although  all  bivalent,  do  not  form  a  coherent  family. 
Glucinum  and  magnesium  resemble  zinc  and  cadmium,  and  differ 
from  the  calcium  family,  in  that  the  sulphates  are  soluble,  the 
hydroxides  easily  lose  water  leaving  the  oxides,  and  the  metals  are 
not  rapidly  rusted  in  the  air  and  do  not  easily  displace  hydrogen 
from  water.  They  resemble  the  calcium  family,  and  differ  from 
zinc  and  cadmium,  in  that  the  sulphides  are  hydrolyzed  by  water, 
the  oxides  are  not  reduced  by  heating  with  carbon,  complex  cations 
are  not  formed  with  ammonia,  and  the  metals  do  not  enter  into 
complex  anions.  But  glucinum  differs  from  magnesium  and 
resembles  zinc  in  that  its  hydroxide  is  acidic  as  well  as  basic.  This 
is  not  unnatural,  since  in  the  periodic  system  it  lies  between  lithium, 
a  metal,  and  boron,  a  non-metal.  Mercury  is  the  only  member  of 
the  group  that  forms  two  series  of  compounds.  These  are  derived 
from  the  oxides  HgO  and  Hg20.  Mercury  approaches  the  noble 
metals  in  the  ease  with  which  its  oxide  is  decomposed  by  heating, 
and  in  the  position  of  the  free  element  in  the  electromotive  series. 

The  vapor  densities  of  zinc,  cadmium,  and  mercury  show  the 
vapors  of  these  three  metals  to  be  monatomic. 

GLUCINUM  Gl. 

Glucinum  (or  beryllium)  is  bivalent  in  all  its  compounds.  Its 
oxide  and  hydroxide  are  basic,  and  are  also  feebly  acidic  towards 
active  bases  (see  Zinc  hydroxide).  The  element  derives  its  name 
from  the  sweet  taste  of  its  salts  (Gk.  yA.v*v?;  sweet). 

427 


428  COLLEGE   CHEMISTRY 

Glucinum  occurs  in  beryl,  a  metasilicate  of  glucinum  and  alumin- 
ium Gl3Al2(SiO3)6.  Beryls,  tinted  green  by  the  presence  of  a  little 
silicate  of  chromium,  are  known  as  emeralds.  The  metal,  obtained 
by  electrolysis  of  the  easily  fusible  double  fluoride  G1F2,  2KF;  burns 
when  heated  in  the  air.  It  displaces  hydrogen  from  dilute  acids, 
and  when  heated,  from  caustic  potash:  Gl  +  2KOH— >K2G1O2 -f  B^. 

MAGNESIUM  Mg. 

Chemical  Relations  of  the  Element.  —  Magnesium  is  bivalent 
in  all  its  compounds.  The  oxide  and  hydroxide  are  basic  exclu- 
sively. The  element  does  not  enter  into  complex  cations  or  anions. 

Occurrence.  —  Magnesium  carbonate  occurs  alone  as  magnesite, 
and  in  a  double  salt  with  calcium  carbonate  MgCO3,CaCO3  as  dolo- 
mite. The  sulphate  and  chloride  are  found  as  hydrates  and  as 
constituents  of  double  salts  (see  below)  in  the  Stassfurt  deposits. 
Olivine  is  the  orthosilicate  Mg2SiO4.  Talc  (soapstone)  is  an  acid 
metasilicate  H2Mg3(SiO3)4.  Serpentine  is  a  hydrated  disilicate, 
[Mg,Fe]3,Si2O7,2H2O,  as  is  also  meerschaum.  Asbestos  is  an  an- 
hydrous silicate.  The  element  derives  its  name  from  Magnesia,  a 
town  in  Asia  Minor. 

The  Metal.  —  Magnesium  is  manufactured  by  electrolysis  of 
dehydrated  and  fused  carnallite  MgCl2,KCl,6H2O.  The  iron  crucible 
in  which  the  material  is  melted  forms  the  cathode,  and  a  rod  of  car- 
bon the  anode.  The  metal  is  silver-white,  and  when  heated  can  be 
pressed  into  wire  and  rolled  into  ribbon. 

Chemically  the  metal  is  less  active  than  are  the  metals  of  the  alka- 
line earths.  It  slowly  becomes  coated  with  a  layer  of  the  carbonate. 
It  displaces  hydrogen  easily  from  boiling  water  and,  of  course,  from 
cold,  dilute  acids.  Magnesium  burns  in  air  with  a  white  light 
(p.  306,  footnote).  The  ash  contains  the  nitride  Mg3N2,  as  well  as 
the  oxide. 

Powdered  magnesium  is  used  in  pyrotechny  and,  with  potassium 
chlorate  (10  : 17),  in  making  flashlight  powder  for  use  in  photography. 

Magnesium  Chloride  MgCl^SH^O.  —  This  highly  deliquescent 
salt  occurs  in  salt  deposits,  alone,  and  as  carnallite  MgCl2,  KC1,  6H20. 
The  latter  is  an  important  source  of  potassium  chloride  (p.  362),  and 


429 

almost  all  the  magnesium  chloride  combined  with  it  is  thrown  away. 
When  the  hexahydrate  is  heated,  a  part  of  the  chloride  is  hydrolyzed, 
some  magnesium  oxide  remaining,  and  some  hydrogen  chloride  being 
given  off.  Sea-water  cannot  be  used  in  ships'  boilers  because  of 
the  hydrochloric  acid  thus  liberated  by  the  magnesium  chloride 
which  the  water  contains.  Anhydrous  magnesium  chloride  MgCl2  is 
obtained  by  heating  the  double  chloride  MgCl2,NH4Cl,6H2O,  for  this 
salt  can  be  dehydrated  without  hydrolysis  of  the  chloride.  The 
ammonium  chloride  is  volatilized  (p.  283). 

The  Oxide  and  Hydroxide.  —  Magnesium  oxide  MgO  is  made 
by  heating  the  carbonate,  and  is  known  as  "  calcined  magnesia."  It 
is  a  white,  highly  infusible  powder,  and  is  used  for  lining  electric 
furnaces  and  making  crucibles.  It  combines  slowly  with  water  to 
form  the  hydroxide  Mg(OH)2. 

The  hydroxide  is  found  in  nature  as  brucite.  It  is  also  precipi- 
tated from  solutions  of  magnesium  salts  by  alkalies.  It  is  very 
slightly  soluble  in  water.  The  solution  has  a  faint  alkaline  reaction. 

Magnesium  hydroxide  is  not  precipitated  by  ammonium  hydroxide 
when  ammonium  salts  are  present  also.  The  ammonium  salts,  being 
highly  ionized  and  giving  a  high  concentration  of  ammonium-ion 
NH4*,  repress  the  ionization  of  the  feebly  ionized  ammonium  hydrox- 
ide, and  so  reduce  the  concentration  of  hydroxide-ion  which  it  fur- 
nishes. With  the  ordinary  concentration  of  Mg",  therefore,  the 
amount  of  hydroxide-ion  existing  in  presence  of  excess  of  a  salt  of 
ammonium  is  too  small  to  bring  the  solubility  product  [Mg**]  X  [OH']2 
up  to  the  value  required  for  precipitation  (cf.  p.  394).  Conversely, 
magnesium  hydroxide  interacts  with  solutions  of  ammonium  salts 
and  passes  into  solution: 

Mg(OH)2(solid)  <=>Mg(OH)2  (dslvd)±=;  Mg"+  20H' 
2NH4C1  ±=>2W   +  2NH4' 

In  presence  of  excess  of  ammonium  chloride,  the  OH'  combines  with 
NH4"  to  form  molecular  ammonium  hydroxide,  and  the  equilibria  in 
the  upper  line  are  displaced  forwards  to  generate  a  further  supply  of 
the  former.  With  sufficiently  great  concentration  of  the  ammonium 
chloride,  all  the  magnesium  hydroxide  may  thus  dissolve.  The 
whole  case  is  analogous  to  the  interaction  of  acids  with  insoluble 
salts  (p.  397). 


430  COLLEGE    CHEMISTRY 

Other  Salts  of  Magnesium.  —  The  normal  carbonate  MgCO3  is 
found  in  nature.  Only  hydrated  basic  carbonates  are  formed  by 
precipitation,  and  their  composition  varies  with  the  conditions. 
The  carbonate  manufactured  in  large  amounts  and  sold  as  magnesia 
alba  is  approximately  Mg4(OH)2(CO3)3.3H20. 

The  common  heptahydrate  of  magnesium  sulphate  MgSO4,7H2O 
crystallizes  from  cold  water  in  rhombic  prisms,  and  is  called  Epsom 
salts.  The  heptahydrate  is  efflorescent.  The  monohydrate  which 
remains,  and  is  found  also  in  the  salt  layers  as  kieserite  MgS04,H2O, 
has  a  very  low  aqueous  tension,  and  is  not  rapidly  dehydrated  except 
above  200°.  Magnesium  sulphate  is  used  in  the  manufacture  of 
sodium  and  potassium  sulphates,  and  is  employed  also  for  "  loading  " 
cotton  goods,  and  as  a  purgative. 

The  sulphide  MgS  may  be  formed  by  heating  the  metal  with  sul- 
phur. It  is  insoluble  in  water,  but  is  decomposed  and  gives,  finally, 
hydrogen  sulphide  and  magnesium  hydroxide: 

MgS  +  2H20  <=»  Mg(OH)2  |  +  H2S. 

The  only  phosphate  of  importance  is  ammonium-magnesium  ortho- 
phosphate  NH4MgPO4,6H20,  which  appears  as  a  crystalline  pre- 
cipitate when  sodium  phosphate  and  ammonium  hydroxide  (and 
chloride,  p.  429)  are  mixed  with  a  solution  of  a  magnesium  salt. 

Analytical   Reactions   of  Magnesium    Compounds.  —  The 

magnesium  ion  is  colorless  and  bivalent.  Soluble  carbonates  pre- 
cipitate basic  carbonates  of  magnesium,  but  not  when  ammonium 
salts  are  present.  The  latter  limitation  distinguishes  compounds  of 
magnesium  from  those  of  the  calcium  family.  Potassium  hydroxide 
precipitates  the  hydroxide  of  magnesium,  except  when  salts  of 
ammonium  are  present.  The  mixed  phosphate  of  ammonium  and 
magnesium,  in  presence  of  ammonium  hydroxide,  is  the  least  soluble 
salt. 

ZINC  Zn. 

Chemical  Relations  of  the  Element.  —  Zinc  is  bivalent  in  all 
its  compounds.  Of  these  there  are  two  sets,  —  the  more  numerous 
and  important  one,  in  which  zinc  is  the  positive  radical  (Zn.S04, 
Zn.Cl2,  etc.),  and  a  less  numerous  set,  the  zincates,  in  which  zinc  is  in 
the  negative  radical  (Na2.Zn02,  etc.).  Both  sets  of  salts  are  hydro- 


431 

lyzed  by  water,  as  the  hydroxide  is  feeble  whether  it  is  considered  as 
an  acid  or  as  a  base.  The  element  also  enters  into  complex  cations 
and  anions.  The  salts  are  all  poisonous. 

Occurrence  and  Extraction  from  the  Ores.  —  The  chief 
sources  of  zinc  are  calamine  or  smithsonite  ZnC03,  zinc-blende  (Ger. 
blenden,  to  dazzle)  or  sphalerite  ZnS,  franklinite  Zn(Fe02)2,  and 
zincite  ZnO. 

The  ores  are  first  converted  into  oxide,  —  the  carbonate  by 
ignition,  and  the  sulphide  by  roasting.  The  sulphur  dioxide  is  used 
to  make  sulphuric  acid.  A  mixture  of  the  oxide  with  coal  is  then 
distilled  in  earthenware  retorts  at  1300-1400°,  the  zinc  condensing 
in  earthenware  receivers,  while  carbon  monoxide  burns  at  a  small 
opening: 

2ZnS  +  302  -»  2ZnO  +  2SO2, 
ZnO  +  C  ->  CO  +  Zn. 

At  first  zinc  dust  (a  mixture  of  zinc  and  zinc  oxide)  collects  in  the 
receiver,  and  afterwards  liquid  zinc.  The  product,  which  is  cast  in 
blocks,  is  called  spelter. 

Properties  and  Uses  of  the  Metal.  —  Zinc  is  a  bluish-white 
crystalline  metal.  When  cold  it  is  brittle,  but  at  120-150°  it  can  be 
rolled  into  sheets  between  heated  rollers  and  then  retains  its  pliability 
when  cold.  At  200-300°  the  metal  becomes  once  more  brittle,  at 
433°  it  melts,  and  at  920°  it  boils.  The  vapor  at  1740°  is  monatomic. 

The  metal  burns  in  air  with  a  greenish  flame,  giving  zinc  oxide.  In 
cold,  moist  air  it  is  oxidized,  and  becomes  covered  with  a  firmly 
adhering  layer  of  basic  carbonate  which  protects  it  from  further 
action.  The  metal  displaces  hydrogen  from  dilute  acids.  Zinc  also 
attacks  boiling  alkalies,  giving  the  soluble  zincate  (see  below): 
2KOH  +  Zn  -*  K2Zn02  +  H2. 

Sheet  zinc,  in  consequence  of  its  lightness  (sp.  gr.  7),  is  used  in 
preference  to  lead  (sp.  gr.  11.5)  for  roofs,  gutters,  and  architectural 
ornaments.  Galvanized  iron  is  made  by  dipping  cleaned  sheet  iron 
in  molten  zinc.  The  latter,  being  more  active  (p.  245),  is  rusted 
instead  of  the  iron.  Zinc  is  used  also  in  batteries  and  for  making 
alloys  (p.  352).  It  mixes  in  all  proportions  with  tin,  copper,  and 
antimony. 


432  COLLEGE   CHEMISTRY 

Zinc  Chloride  ZnCl2.  —  This  salt  is  usually  manufactured  by 
treating  zinc  with  excess  of  hydrochloric  acid,  evaporating  the 
solution  to  dryness,  and  fusing  the  residue.  When  hydrochloric 
acid  is  thus  present,  the  chloride  ZnCl2  is  obtained.  Evaporation  of 
the  pure  aqueous  solution,  which  is  acid  in  reaction,  results  in  con- 
siderable hydrolysis  and  formation  of  much  of  the  basic  -chloride 
Zn2OCl2: 

ZnCl2  +  H2O  <=±  HC1  +  Zn(OH)Cl,  (1) 

2Zn(OH)Cl  -»  Zn2OCl2  +  H,O.  (2) 

The  salt  is  used  in  solid  form  as  a  caustic  and,  by  injection  of  a  solu- 
tion into  wood  (e.g.,  railway  sleepers),  as  a  poison  to  prevent  the 
growth  of  organisms  which  promote  decay.  In  both  cases  the  salt 
combines  with  albumins,  forming  solid  products. 

r 

tZinc  Oxide  ^and  Hydroxide  and  the  Zincates.  —  The  oxide 

lO  is  obtained  as  a  white  powder  by  burning  zinc  or  by  heating  the 
precipitated  basic  carbonates.  It  turns  yellow  when  heated, 
recovering  its  whiteness  when  cold,  in  the  same  way  that  mercuric 
oxide  is  brown  whilst  hot  and  bright  red  when  cold.  It  is  employed 
in  making  a  paint  —  zinc-white  or  Chinese  white  — \  which  is  not 
darkened  by  hydrogen  sulphide. 

The  hydroxide  Zn(OH)2  appears  as  a  white,  flocculent  solid  when 
alkalies  are  added  to  solutions  of  zinc  salts.  It  interacts  as  a  basic 
hydroxide  with  acids,  giving  salts  of  zinc : 

Zn(OH)2  -f  H2SO4  <=>  Zn.SO4  +  2H20. 

It  also  interacts  with  excess  of  the  alkali  employed  to  precipitate  it, 
giving  a  soluble  zincate,  such  as  potassium  zincate  K2Zn02: 

H2ZnO  2  |  +  2KOH  <=*  K2.Zn02  +  2H,O. 
Zinc  hydroxide  is  ionized  both  as  an  acid  and  as  a  base: 


2H'  +  ZnO;"  <=±  Zn(OH)2  (dslvd)  <=±  Zn"  +  2QH' 
Zn(OH)2  (solid) 

The  ionization  as  an  acid  is  less  than  that  as  a  base,  but  both  are 
small.  Addition  of  an  acid  like  sulphuric  acid,  however,  furnishes 
hydrogen-ion;  the  hydroxyl  ions  combine  with  this  to  form  water, 
and  all  the  equilibria  are  displaced  to  the  right.  With  a  base,  on  the 


ZINC  433 

I 

other  hand,  the  hydrogen-ion  is  removed  and  the  basic  ionization 
simultaneously  repressed,  so  that  the  equilibria  are  displaced  to  the 
left.  t 

Zinc  hydroxide  interacts  with  ammonium  hydroxide,  giving  the 
soluble  ammonia-zinc  hydroxide  Zn(NH3)4.(OH)2.  The  case  is  like 
those  of  copper  (p.  414)  and  silver  hydroxides  (p.  420). 

Compounds  of  zinc,  when  heated  in  the  Bunsen  flame  with  a  salt 
of  cobalt,  gives  a  zincate  of  cobalt  (Rinmann's  green)  CoZn02. 

Other  Salts  of  Zinc.  —  The  normal  zinc  carbonate  ZnC03  may 
be  precipitated  by  means  of  sodium  bicarbonate,  but  normal  carbo- 
nate of  sodium  gives  basic  carbonates,  such  as  Zn2(OH)2C03: 

2ZnS04  +  2Na2C03  +  H20  ->  Zn2(OH)2C03  +  2Na2S04  +  C02. 


Zinc  sulphate  ZnS04  is  'formed  when  zinc-blende  is  roasted.  It 
gives  rhombic  crystals  of  the  hydrate  ZnSO4,  7H20.  This,  and  the 
corresponding  compounds  of  magnesium  MgSO4,7H2O,  of  iron  FeSO4, 
7H20,  and  of  other  bivalent  metals  are  known  as  vitriols.  The  zinc 
salt  is  white  vitriol.  It  is  used  in  cotton-printing  and  as  an  eye- 
wash (J  per  cent  solution).  The  sulphate  gives  double  salts,  such  as 
potassium-zinc  sulphate  ZnSO4,K2SO4,6H2O  (cf.  p.  231). 

Zinc  sulphide  ZnS  is  more  soluble  in  water  than  is  sulphide  of 
copper,  and  hence  it  interacts  with  excess  of  strong  acids,  and  passes 
into  solution.  It  is  not  soluble  enough,  however,  to  be  much  affected 
by  weak  acids  like  acetic  acid  (cf.  p.  398).  Zinc  sulphide  is  thus 
capable  of  being  precipitated  when  acetic  acid  is  present,  or  when 
hydrogen  sulphide  is  led  into  a  solution  of  the  acetate  of  zinc  : 

Zn(C2H302)2  +  H2S  <=>  ZnS  |  +  2HC2H302. 

But  when  an  active  acid  is  present,  or  is  formed,  the^jlpjude  is 
precipitated  incompletely  or  not  at  all,  the  action  being  reversible: 

ZnSO4  +  H2S  <=»  ZnS  +  H2S04. 

There  are  thus  two  ways  of  obtaining  the  sulphide  by  precipita- 
tion. A  soluble  sulphide  causes  it  to  be  thrown  down  completely, 
because  no  acid  is  liberated  in  the  action: 

ZnCl2  +  (NH4)2S  <=±  ZnS  J  4-  2NH4C1. 

The  other  method  is  to  add  sodium  acetate  to  the  solution  of  the  salt, 
and  then  lead  in  hydrogen  sulphide.     The  acid  which  is  liberated  by 


434  COLLEGE   CHEMISTRY 

the  action  upon  the  salt  interacts  with  the  sodium  acetate,  giving  a 
neutral  salt  of  sodium  and  acetic  acid,  and  the  zinc  sulphide  is  not 
affected  by  the  latter  (cf.  p.  399). 

Analytical  Reactions  of  Zinc  Salts.  —  Zinc  sulphide  is  pre- 
cipitated by  the  addition  of  ammonium  sulphide  to  solutions  of  zinc 
salts  and  of  zincates.  Sodium  hydroxide  gives  the  insoluble  hydrox- 
ide, which,  however,  interacts  with  excess  of  the  alkali,  giving  the 
soluble  zincate  of  sodium.  Compounds  of  zinc,  when  heated  on 
charcoal  •Cobalt  nitrate,  give  Rinmann's  green  (p.  433). 

CADMIUM  Cd. 

Chemical  Relations  of  the  Element.  —  This  element  is  biva- 
lent in  all  its  compounds.  Its  oxide  and  hydroxide  are  basic  exclu- 
sively, and  the  salts  are  not  hydrolyzed  by  water.  It  enters  into 
complex  compounds  having  the  ions  Cd  (NH3)  4"  and  Cd  (CN)  /'.  Its 
resemblances  to,  and  differences  from  zinc  are  especially  note- 
worthy. | 

The  Metal.  —  Aside  from  the  rare  mineral  greenockite  CdS, 
cadmium  is  found  only  in  small  amounts,  as  carbonate  and  sulphide, 
in  the  corresponding  ores  of  zinc.  During  the  reduction,  being  more 
volatile  than  zinc,  it  distils  over  first  (b.-p.  770°).  The  metal  is 
white,  and  is  more  malleable  than  zinc.  It  displaces  hydrogen  from 
dilute  acids  (cf.  p.  245). 

Compounds  of  Cadmium.  —  The  chloride  CdCl2,2H20  is  efflores- 
cent and  is  not  hydrolyzed  during  dehydration  or  in  solution.  Zinc 
chloride  (p.  432)  is  deliquescent  and  is  easily  hydrolyzed. 

The  hydroxide  Cd(OH)2  is  made  by  precipitation,  and  interacts 
with  acids  (as  a  basic  hydroxide),  but  not  at  all  with  bases.  It  dis- 
solves in  ammonium  hydroxide,  however,  forming  Cd(NH3)4.(OH)2. 
The  oxide  CdO  is  a  brown  powder,  obtained  by  heating  the  hydrox- 
ide^ carbonate,  or  nitrate,  or  by  burning  the  metal. 

The  sulphate  crystallizes  from  solution  as  3CdS04,8H2O.  Soluble 
carbonates  throw  down  the  normal  carbonate  of  cadmium  CdC03. 

Hydrogen  sulphide  precipitates  the  yellow  sulphide  CdS  even  from 
acid  solutions  of  the  salts.  The  substance  is  used  as  a  pigment.  The 


j   copper, 


MERCURY  485 

sulphide  of  cadmium,  however,  is  less  insoluble  in  water  (cf.  p.  397) 
than  are  the  sulphides  of  copper  and  mercury,  and  therefore  cannot 
be  precipitated  from  a  very  strongly  acid  solution. 

Analytical  Reactions  of  Cadmium  Compounds*  —  The  cad- 
mium ion  Cd"  is  bivalent  and  colorless.  The  yellow  cadmium  sul- 
phide is  precipitated  by  hydrogen  sulphide,  even  from  acid  solutions 
of  the  salts.  The  white,  insoluble  hydroxide  is  not  soluble  in  sodium 
hydroxide. 

MERCURY  Hg. 

Chemical  Relations  of  the  Element.  —  Like  HP1F,  this 
element  enters  into  two  series  of  compounds,  the  mercurous  Hg1  and 
the  mercuric  Hg11.  The  mercurous  halides,  like  the  cuprous  halides 
(and  the  argentic  halides),  are  insoluble  in  water  and  are  decom- 
posed by  light.  Both  of  the  oxides,  Hg2O  and  HgO,  are  basic  exclu- 
sively, but  in  a  feeble  degree.  The  hydroxides,  like  silver  hydroxide, 
are  not  stable,  and  lose  water,  giving  the  oxides.  The  salts  of  both 
sets  are  markedly  hydrolyzed  by  water,  and  basic  salts  are  therefore 
common.  No  carbonate  is  known.  Mercury  enters  into  the  anions 
of  a  number  of  complex  salts,  such  as  HgCl/',  HgI4",  'Hg(CN)/', 
etc.  It  forms  a  class  of  mercury-ammonia  compounds,  like  Hg"NH2Cl, 
all  of  which  are  insoluble. 

The  mercury  salts  of  volatile  acids,  like  the  corresponding  salts  of 
ammonium  (p.  283),  can  all  be  volatilized  completely.  Mercury 
vapor  and  all  mercury  compounds  are  poisonous,  the  soluble  ones 
more  markedly  so  than  the  insoluble  ones. 

Occurrence  and  Isolation  of  the  Metal.  —  Mercury  occurs 
native  and  to  a  larger  extent  as  red,  crystalline  cinnabar,  mercuric 
sulphide  HgS.  The  chief  mines  are  in  Spain,  California,  arid 
Austria. 

The  liberation  of  the  metal  is  easy,  because,  when  roasted,  the 
sulphide  is  decomposed,  and  the  sulphur  forms  sulphur  dioxide. 
The  mercury  does  not  unite  with  oxygen,  for  the  oxide  decomposes 
(p.  7)  at  400-600°: 


In  some  places  the  ore  is  spread  on  perforated  brick  shelves  in  a  verti- 
cal furnace,  and  the  gases  pass  through  tortuous  flues  in  which  the 
vapor  of  the  metal  condenses. 


436  COLLEGE   CHEMISTRY 

Physical  Properties.  —  -  Mercury  or  quicksilver  (N.L.  hydrargy- 
rum, from  Gk.  v&op,  water,  and  apyupos,  silver)  is  a  silver-white 
liquid.  At  —  39.5°  it  freezes,  and  at  357°  it  boils. 

On  account  of  its  high  specific  gravity  (13.6,  at  0°)  and  low  vapor 
tension,  the  metal  is  employed  for  filling  barometers.  Its  uniform 
expansion  favors  its  use  in  thermometers.  The  tendency  to  form 
amalgams,  which  it  exhibits  towards  all  the  familiar  metals  with 
the  exception  of  iron  and  platinum  (the  latter,  however,  is  "  wet  "  by 
it),  is  taken  advantage  of  in  various  ways  (cf.  pp.  283,  423). 

Chemical  Properties.  —  When  kept  at  a  temperature  near  to  its 
boiling-point,  mercury  combines  slowly  with  oxygen.  Mercury  does 
not  displace  hydrogen  from  dilute  acids  (p.  245),  but  with  oxidizing 
acids  like  nitric  acid  and  hot  concentrated  sulphuric  acid,  the  nitrates 
and  sulphate  (mercuric)  are  formed.  With  excess  of  mercury, 
mercurous  nitrate,  and  with  excess  of  the  hot  acid,  mercuric  nitrate, 
are  produced.  When  mercury  is  divided  into  minute  droplets,  with 
relatively  large  surface,  it  is  used  in  medicine  ("  blue  pills  "),  and 
shows  an  activity  which  is  entirely  wanting  in  larger  masses. 

TJie  Halides  of  Mercury.  —  Mercurous  chloride  HgCl  (calomel) 
is  obtained  as  a  white  powder  by  precipitation  from  solutions  of 
mercurous  salts.  It  is  made  by  subliming  mercuric  chloride  with 


or  more  usually  by  subliming  a  mixture  of  mercuric  sulphate,  made 
as  described  above,  with  mercury  and  common  salt.  It  is  deposited 
on  the  cool  part  of  the  vessel  as  a  fibrous  crystalline  mass,  or,  when 
the  vapor  is  led  into  a  large  chamber,  as  a  fine  powder.  It  is  slowly 
affected  by  light  just  as  silver  chloride  is.  Here,  however,  the  chlo- 
rine which  is  released  combines  with  another  molecule  of  the  salt  to 
form  mercuric  chloride.  The  substance  is  used  in  medicine  on 
account  of  its  tendency  to  stimulate  all  organs  producing  secretions. 
By  direct  union  with  chlorine,  mercuric  chloride  HgCl2  (corrosive 
sublimate)  is  formed.  It  is  usually  manufactured  by  subliming 
mercuric  sulphate  with  common  salt,  and  crystallizes  in  white, 
rhombic  prisms.  It  melts  at  265°  and  boils  at  307°.  The  solubility 
at  20°  is  7.4  :  100  Aq.  The  aqueous  solution  is  slightly  acid  in 
reaction.  The  salt  is  easily  reduced  to  mercurous  chloride.  When 


MERCURY  437 

excess  of  stannous  chloride  is  added  to  the  solution,  the  white  pre- 
cipitate of  calomel,  first  formed,  passes  into  a  heavy  gray  precipitate 
of  finely  divided  mercury: 

2HgCl2  +  SnCl2  ->  SnCl4  +  2HgCl, 
2HgCl   +  SnCl2  ->  SnCl4  +  2Hg. 

Corrosive  sublimate,  when  taken  internally,  is  extremely  poisonous. 
A  very  dilute  solution  is  used  in  surgery  to  destroy  lower  organisms 
and  thus  prevent  infection  of  wounds.  Mercuric  chloride  acts  also 
as  a  preservative  of  zoological  materials,  forming  insoluble  compounds 
with  albumins,  and  preventing  decay.  For  the  same  reason,  albumin 
(white  of  an  egg)  is  given  as  an  antidote  in  cases  of  sublimate 
poisoning. 

Mercurous  iodide  Hgl  is  formed  by  rubbing  iodine  with  excess  of 
mercury.  It  also  appears  as  a  greenish-yellow  precipitate  when 
potassium  iodide  is  added  to  a  solution  of  a  mercurous  salt.  The 
compound  decomposes  spontaneously  into  mercury  and  mercuric 

iodide:  ~TT  T      TT     ,  TT  T 

2HgI  <=>  Hg  +  HgI2. 

Mercuric  iodide  HgI2  is  obtained  by  direct  union  of  mercury  with 
excess  of  iodine,  or  by  addition  of  potassium  iodide  to  a  solution  of 
a  mercuric  salt.  It  is  a  scarlet  powder,  insoluble  in  water,  but 
soluble  in  alcohol  and  ether.  It  interacts  with  excess  of  potassium 
iodide,  forming  the  soluble,  colorless  potassium  mercuri-iodide  K2.HgI4 
with  which  many  precipitants  fail  to  give  mercury  compounds. 

TJie  Oxides.  —  When  bases  (excepting  ammonium  hydroxide, 
see  p.  438)  are  added  to  solutions  of  mercurous  salts,  the  greenish- 
black  mercurous  oxide  Hg2O  is  thrown  down.  The  hydroxide  is 
doubtless  formed  transitorily  and  then  loses  water  (cf.  Silver  oxide, 
p.  420).  Under  the  influence  of  light  or  gentle  heat  (100°),  this 
oxide  resolves  itself  into  mercuric  oxide  and  mercury. 

Mercuric  oxide  HgO  is  formed  as  a  red,  crystalline  powder,  when 
mercury  is  heated  in  air  near  to  357°,  but  is  usually  made  by  decom- 
posing the  nitrate.  Commercial  specimens,  incompletely  decom- 
posed, thus  give  some  nitrogen  peroxide  when  heated.  It  is  formed 
also  as  a  yellow  powder  by  adding  bases  to  solutions  of  mercuric 
salts. 


438  COLLEGE   CHEMISTRY 

Other  Salts  of   Mercury.  —  Mercurous  nitrate  HgNO3,H2O  is 

formed  by  the  action  of  cold,  diluted  nitric  acid  upon  excess  of 
mercury.  It  is  hydrolyzed,  slowly  by  cold,  and  rapidly  by  warm 
water,  giving  a  basic  nitrate : 

2HgNO,  +  H2O<=>HNO3  +  Hg2(OH)NO3J. 

On  this  account  a  clear  solution  can  be  made  only  when  some  nitric 
acid  is  added.  Free  mercury  is  also  kept  in  the  solution  to  reduce 
mercuric  nitrate,  which  is  formed  by  atmospheric  oxidation: 

Hg(N03)2  +  Hg->2HgN03    or    Hg"  +  Hg -»  2Hg\ 

Mercuric  nitrate  Hg(N03)2,8H2O  is  produced  by  using  excess  of 
warm,  concentrated  nitric  acid  with  mercury.  The  aqueous 
solution  is  strongly  acid,  and  deposits  a  yellowish,  crystalline,  basic 
nitrate  Hg3(OH)2O(N03)2.  The  hydrolysis  is  reversed  by  adding 
nitric  acid. 

Mercurous  sulphide  Hg2S  is  formed  by  precipitation  from  mer- 
curous  salts,  but  decomposes  into  mercury  and  mercuric  sulphide. 

Crystallized  mercuric  sulphide  HgS  occurs  as  cinnabar,  and  is  red. 
When  formed  by  precipitation  with  hydrogen  sulphide,  or  by  rubbing 
together  mercury  and  sulphur,  it  is  black  and  amorphous.  By 
sublimation,  in  the  course  of  which  it  dissociates,  the  black  form 
gives  the  red,  crystalline  one. 

The  black  and  the  red  varieties  do  not  interact  with  concentrated 
acids,  or  even  with  boiling  nitric  acid,  which  oxidizes  most  sulphides 
readily.  They  are,  therefore,  still  less  soluble  than  is  cupric  sulphide 
(pp.  397-398).  They  are  attacked,  however,  by  aqua  regia.  The 
red  form  of  the  sulphide  is  used  in  making  paint  (vermilion). 

Mercuric  fulminate  Hg(ONC)2  is  obtained  as  a  white  precipitate 
when  mercury  is  treated  with  nitric  acid,  and  alcohol  is  added  to  the 
solution.  It  decomposes  suddenly  when  struck,  and  is  used  in 
making  percussion  caps. 

Mercury -ammonia  Compounds. — When  ammonium  hydrox- 
ide is  added  to  a  solution  of  a  mercuric  salt,  a  white  substance,  of  a 
type  which  we  have  not  previously  encountered,  is  thrown  down. 
Mercuric  chloride  gives  HgNH2Cl,  commonly  called  "  infusible 
white  precipitate,"  and  often  mercuri-ammonium  chloride: 

HgCl2  +  2NH3  ->  HgNH2Cl  +  NH4C1. 


RECOGNITION  OF  CATIONS  IN  QUALITATIVE  ANALYSIS     439 

Mercuric  nitrate  gives  Hg2NNO3,  sometimes  called  dimercuri- 
ammonium  nitrate : 

2Hg(N03)2  +  4NH3-»Hg2NN03  +  3NH4NO3. 

The  proper  classification  of  these,  and  the  other  similar  substances, 
is  beyond  the  range  of  this  book. 

When  calomel  is  treated  with  ammonium  hydroxide,  it  turns  into 
a  black,  insoluble  body.  This  appears  to  be  a  mixture  of  free 
mercury,  to  which  it  owes  its  dark  color,  and  "  infusible  white 
precipitate,"  Hg  +  HgNH2Cl.  To  this  reaction  calomel  owes  its 
name  (Gk.  KaXo^e\as,  beautiful  black).  Mercurous  nitrate  gives  a 
black,  insoluble  mixture,  2Hg  +  Hg2NN03. 

Analytical  Reactions  of  Mercury  Compounds.  —  The  two 

ionic  forms  of  the  element,  mercurous-ion  Hg'  and  mercuric-ion 
Hg",  are  both  colorless.  Their  chemical  behavior  is  entirely 
different.  Both  give  the  black  sulphide  HgS,  which  is  insoluble  in 
acids  and  other  solvents  of  mercury  salts.  Mercurous-ion  gives  the 
insoluble,  white  chloride,  the  black  oxide,  and  a  black  mixture  with 
ammonium  hydroxide.  Mercuric-ion  gives  a  soluble  chloride,  a 
yellow,  insoluble  oxide,  and  a  white  precipitate  with  ammonium 
hydroxide.  The  behavior  with  stannous  chloride  (p.  437)  is  char- 
acteristic. With  potassium  iodide  the  two  ions  behave  differently 
(p.  437).  The  more  active  metals  displace  mercury  from  all  com- 
pounds. Copper  is  used  as  the  displacing  metal  because  the  mer- 
cury is  easily  seen  on  its  surface. 

Salts  of  mercury  are  volatile.  When  heated  in  a  tube  with 
sodium  carbonate,  they  give  a  sublimate  of  metallic  mercury. 

THE  RECOGNITION  OF  CATIONS  IN  QUALITATIVE  ANALYSIS. 

"  Wet-way  "  analysis  consists  in  recognizing  the  various  positive 
and  negative  ions  present  in  a  solution  (p.  232).  In  discussing 
hydrogen  sulphide  (p.  254),  it  was  stated  that  the  sulphides  might 
be  divided  into  three  classes  according  to  their  behavior  towards 
water  and  acids.  Now  these  differences  furnish  us  with  a  basis  for 
distinguishing  the  cations  present  in  a  solution. 

The  following  plan,  taken  in  conjunction  with  the  statements  in 
the  context,  shows  how  a  single  cation  may  be  identified,  and  how, 
when  several  cations  are  present,  a  separation  preparatory  to  identifi- 


440  COLLEGE   CHEMISTRY 

cation  may  be  effected.  What  will  be  said  applies  only  to  the  case 
of  a  solution  containing  salts  like  the  chlorides,  nitrates,  or  sulphates 
of  one  or  more  cations,  and  leaves  the  oxalates  (p.  399),  phosphates, 
cyanides,  and  some  other  salts,  out  of  consideration. 

Group  1.  —  Add,  first,  hydrochloric  acid,  to  find  out  whether  cations 
giving  insoluble  chlorides  are  present.  Argentic,  mercurous,  and 
plumbic  salts  give  the  white  AgCl,  HgCl,  and  PbCl2,  respectively 
(cf.  p.  422).  Filtration  eliminates  the  precipitate,  if  there  is  any. 

Group  2.  —  A  free,  active  acid  being  now  present,  hydrogen  sulphide 
is  led  into  the  solution.  The  sulphides  insoluble  in  active  acids, 
namely,  HgS,  CuS,  PbS,  Bi^Sg,  CdS,  As2S3,  Sb2S3,  SnS,  SnS2  are  there- 
fore thrown  down.  The  first  four  are  black  or  brown,  the  next  two 
and  the  last  are  yellow,  and  the  remaining  two  are  orange  and 
brown  respectively.  A  dark-colored  substance  will  naturally  obscure 
one  of  lighter  color,  if  more  than  one  is  present.  Filtration  again 
eliminates  the  precipitate. 

This  group  is  easily  subdivided.  Any  or  all  of  the  last  four  sul- 
phides will  pass  into  solution  when  warmed  with  yellow  ammonium 
sulphide,  for  they  give  soluble  complex  sulphides  (q.v.).  The  first 
five  sulphides,  or  any  of  them,  will  be  unaffected.  On  the  other 
hand,  these  five  sulphides,  with  the  exception  of  HgS,  will  interact 
with  hot  nitric  acid  (p.  438).  Other  reactions  are  then  used  to 
distinguish  between,  or,  if  there  is  a  mixture,  to  separate,  the  mem- 
bers of  the  sub-groups. 

Group  3.  —  The  solution  (filtrate)  is  now  neutralized  with  ammo- 
nium hydroxide,  and  ammonium  sulphide  is  added.  Some  ammonium 
chloride  is  also  used,  to  prevent  the  precipitation  of  magnesium 
hydroxide  (p.  429),  which,  in  any  event,  would  be  incomplete.  The 
sulphides  which  are  insoluble  in  water,  and  are  not  hydrolyzed  by 
it,  now  appear.  They  are  FeS,  CoS,  NiS,  all  black,  MnS,H2O,  and 
ZnS,  which  are  pink  and  white  respectively.  There  are  precipitated 
also  the  hydroxides  of  chromium  and  of  aluminium,  Cr(OH)3  and 
A1(OH)3,  because  their  sulphides  are  hydrolyzed  by  water. 

Group  4.  —  After  filtration,  ammonium  carbonate  is  added,  and 
precipitates  the  remaining  metals  whose  carbonates  are  insoluble, 
BaCO3,  SrC03>  CaCO3>  with  the  exception  of  magnesium  (p.  430). 

By  addition  of  sodium  phosphate  to  a  portion  of  the  filtrate,  mag- 
nesium, if  present,  now  comes  out  in  the  form  NH4MgP04.  There 
remain  in  solution  only  salts  of  potassium,  sodium,  and  ammonium. 


RECOGNITION  OF   CATIONS  IN  QUALITATIVE  ANALYSIS      441 

Since  only  ammonium  compounds,  and  other  substances  which  can 
be  volatilized  have  been  added,  evaporation  and  ignition  of  the 
residue  leaves  the  salts  of  the  two  metals.  Salts  of  ammonium  must 
be  sought  in  a  fresh  sample  by  the  usual  test  (p.  283). 

Exercises. —  1.  Why  should  we  expect  ammonium  sulphide 
solution  to  precipitate  magnesium  hydroxide,  and  why  does  it  not 
do  so? 

2.  What  volume  of  air  is  required  to  oxidize  one  formula-weight 
of  zinc  sulphide  to  ZnO  and  SO2,  and  what  volume  of  sulphur  dioxide 
is  produced?    Is  the  gaseous  product  more,  or  less  diluted  with 
nitrogen  than  when  pure  sulphur  is  burned,  and  by  how  much? 

3.  Make  equations  showing,  (a)  the  effect  of  heating  zinc  chloride 
with  cobalt  nitrate  (Co(NO3)2)  in  the  Bunsen  flame  (p.  433),  (6)  the 
action  of  hydrogen  sulphide  on  sodium  zincate,  (c)  the  actions  of 
concentrated  nitric  acid  and  of  concentrated  sulphuric  acid  on 
mercury. 

4.  What  kind  of  salts  might  take  the  place  of  sodium  acetate  in 
the  precipitation  of  zinc  sulphide  (p.  433,  foot)?     Give  examples. 

5.  Why  do  none  of  the  salts  of  the  elements  in  this  family  give 
recognizable  effects  with  the  borax  bead? 


CHAPTER  XXXVIII 
ALUMINIUM    AND    THE    METALS    OF    THE   EARTHS 

THE  chief  members  of  the  family  occupying  the  fourth  column  of 
the  periodic  table  are:  boron  (B,  at.  wt.  11),  aluminium  (Al,  at.  wt. 
27.1),  gallium  (Ga,  at.  wt.  70),  indium  (In,  at.  wt.  115),  thallium 
(Tl,  at.  wt.  204.1),  all  on  the  right  side  of  the  column;  and  scandium 
(Sc,  at.  wt.  44.1),  yttrium  (Y,  at.  wt.  89),  lanthanum  (La,  at.  wt. 
138.9),  on  the  left  side.  These  elements  are  all  trivalent. 

The  Rare  Elements  of  this  Family.  —  The  oxide  and  hydrox- 
ide of  boron  are  acidic  (p.  349).  Those  of  aluminium  (A1(OH)3), 
gallium  (Ga(OH)3),  indium  (In(OH)3),  and  thallium  (T1O.OH)  are 
basic,  but  behave  also  as  acids  towards  strong  bases. 

Gallium  and  indium  occur  occasionally  in  zinc-blende,  and  were 
discovered  by  the  use  of  the  spectroscope.  The  former  takes  its 
name  from  the  country  (France)  in  which  the  discovery  was  made, 
and  the  latter  from  two  blue  lines  shown  by  its  spectrum. 

Thallium  is  found  in  some  specimens  of  pyrite  and  blende.  It  was 
discovered  by  Crookes,  by  means  of  the  spectroscope,  in  the  seleni- 
ferous  deposit  from  the  flues  of  a  sulphuric  acid  factory.  It  received 
its  name  from  the  prominent  green  line  in  its  spectrum  (Gk.  0aAAds,  a 
green  twig).  It  gives  two  complete  series  of  compounds.  In  those 
in  which  it  is  trivalent  (thallic  salts),  it  resembles  aluminium  (q.v.). 
Thus,  the  salts  of  this  series  are  more  or  less  hydrolyzed  by  water. 
Univalent  thallium  recalls  both  sodium  and  silver.  Thallous  hydrox- 
ide (T1OH)  is  soluble,  and  gives  a  strongly  alkaline  solution.  The 
chloride  is  insoluble  in  cold  water.  The  solutions  of  the  thallous 
salts  are  neutral.  The  metal  is  displaced  from  its  salts  by  zinc. 

Of  the  elements  on  the  left  side  of  the  column,  scandium,  whose 
existence  and  properties  were  predicted  by  Mendelejeff  (p.  276),  is 
the  best  known.  The  metals  of  the  rare  earths,  of  which  it  is  one,  are 
found  in  rare  minerals  such  as  euxenite,  gadolinite,  orthite,  and 
monazite,  which  occur  in  Sweden,  Greenland,  and  the  United  States. 

442 


ALUMINIUM  443 

Cerium  (Ce,  at.  wt.  140.25),  neodymium  (Nd,  at.  wt.  143.6),  and 
praseodymium  (Pr,  at.  wt.  140.5),  occur  along  with  lanthanum  in 
cerite,  a  silicate  of  these  four  elements.  These  four  are  included 
amongst  the  metals  of  the  rare  earths.  The  compounds  of  many  of 
these  rare  elements  behave  so  much  alike  that  separation  is  difficult. 
It  is  certain,  however,  that  there  are  several  with  atomic  weights 
near  to  that  of  lanthanum  for  which  accommodation  cannot  easily 
be  found  in  the  periodic  table.  Ostwald  has  compared  them  to  a 
group  of  minor  planets  such  as  in  the  solar  system  takes  the  place  of 
one  large  planet. 

ALUMINIUM. 

The  Chemical  Relations  of  the  Element.  —  Aluminium  is 
trivalent  exclusively.  Its  hydroxide,  like  that  of  zinc  (p.  432),  is 
feebly  acidic  as  well  as  basic,  and  hence  the  metal  forms  two  sets  of 
compounds  of  the  types  Na3.AlO3  (sodium  aluminate)  and  A12.(S04)3. 
The  salts  of  both  series  are  more  or  less  hydrolyzed  by  water,  the 
former  very  conspicuously  so.  It  is  worth  noting  that  the  hydrox- 
ides of  the  trivalent  metals,  or  metals  in  the  trivalent  condition,  such 
as  A1(OH)3,  Cr(OH)3,  Fe(OH)3,  are  all  distinctly  less  basic  than  are 
those  of  the  bivalent  metals  such  as  Zn(OH)2,  Cd(OH)2,  Fe(OH)2, 
Mn(OH)2.  Aluminium  does  not  enter  into  complex  anions  or  cations, 
and  is  too  feebly  base-forming  to  give  salts  like  the  carbonate  or 
sulphite. 

Occurrence.  —  Aluminium  is  found  very  plentifully  in  combina- 
tion, coming  next  to  oxygen  and  silicon  in  this  respect.  The  feldspars 
(such  as  KAlSi3O8),  the  micas  (such  as  KAlSiO4),  and  kaolin  (clay 
H2A12  (SiO4)2,H20) ,  are  the  commonest  minerals  containing  it.  Cryo- 
lite is  a  double  fluoride  3NaF,AlF3.  Various  forms  of  the  oxide  and 
hydroxide  are  also  found. 

Preparation  and  Physical  Properties.  —  The  metal  is  now 
made  on  a  large  scale  by  electrolysis  of  the  oxide  (A1203)  dissolved  in 
a  bath  of  molten  cryolite.  The  operation  is  conducted  in  cells,  the 
carbon  linings  of  which  form  the  cathodes.  The  anodes  are  rods  of 
carbon  which  combine  with  the  oxygen  as  it  is  liberated.  The  metal 
sinks  to  the  bottom  of  the  cell  and  is  drawn  off  periodically,  while 
fresh  portions  of  the  oxide  are  added  from  time  to  time.  The  current 


444  COLLEGE   CHEMISTRY 

(E.M.F.  5-6  volts)  maintains  the  temperature  of  the  molten  mate- 
rials, and  causes  the  decomposition. 

The  metal  melts  at  600-700°,  but  is  not  mobile  enough  to  make 
castings.  It  is  exceedingly  light  (sp.  gr.  2.6),  and  in  hardness  and 
tensile  strength  excells  the  other  metals,  with  the  exception  of 
iron  and  copper.  It  has  a  silvery  luster,  and  tarnishes  very  slightly, 
the  firmly  adhering  film  of  oxide  first  formed  protecting  its  surface. 
Although,  comparing  cross-sections,  it  is  not  so  good  a  conductor  of 
electricity  as  is  copper,  yet  weight  for  weight  it  conducts  better.  It 
is  difficult  to  work  on  the  lathe  or  to  polish,  because  it  sticks  to  the 
tools,  but  the  alloy  with  magnesium  (6-30  per  cent)  called  magna- 
lium  has  admirable  qualities  in  these  respects.  Aluminium  bronze 
(5-12  per  cent  aluminium)  is  easily  fusible,  has  a  magnificent  golden 
luster,  and  possesses  mechanical  and  chemical  resistance  exceeding 
that  of  any  other  bronze.  The  metal  and  its  alloys  are  used  for 
making  cameras,  opera-glasses,  cooking  utensils,  and  other  articles 
requiring  lightness  and  strength,  as  well  as  for  the  transmission  of 
electricity.  The  powdered  metal,  mixed  with  oil,  is  used  in  making 
a  silvery  paint. 

Chemical  Properties.  —  The  metal  displaces  hydrogen  from 
hydrochloric  acid  very  easily.  It  displaces  hydrogen  also  from 
boiling  solutions  of  the  alkalies,  forming  aluminates : 

2A1  +  6NaOH  -»  2Na3A103  +  3H2. 

In  consequence  of  its  very  great  affinity  for  oxygen,  aluminium 
displaces  all  the  metals,  save  magnesium,  from  their  oxides.  Thus, 
when  a  mixture  of  aluminium  powder  and  ferric  oxide  is  placed  in  a 
crucible  and  ignited  by  means  of  a  piece  of  burning  magnesium 
ribbon,  aluminium  oxide  and  iron  are  formed: 

Fe,O,  +  2A1  ->  A1203  +  2Fe. 

The  very  high  temperature  (about  3000°)  produced  by  the  action  is 
sufficient  to  melt  both  the  iron  (m.-p.  1550°)  and  the  oxide  of  alumi- 
nium. The  products,  not  being  miscible,  separate  into  two  layers. 
This  very  simple  method  of  making  pure  specimens  of  metals  like 
chromium,  uranium,  and  manganese,  whose  oxides  are  otherwise 
hard  to  reduce,  is  called  by  Goldschmidt,  the  inventor,  alumino- 
thermy.  The  sulphides,  such  as  pyrite,  are  reduced  with  like  vigor 
by  aluminium. 


ALUMINIUM  445 

Aluminium  Chloride  AIC13.  —  If  the  metal  or  the  hydroxide  is 
treated  with  hydrochloric  acid,  and  the  solution  is  allowed  to  evapo- 
rate, the  hydrated  chloride  A1C13,6H2O  is  formed.  When  heated, 
this  hydrate  is  completely  hydrolyzed,  hydrochloric  acid  is  given  off, 
and  only  the  oxide  remains.  The  anhydrous  chloride  A1C13  is  made  by 
passing  dry  chlorine  over  aluminium,  or  by  heating  the  oxide  with 
carbon  in  a  stream  of  chlorine  (cf.  p.  344).  Since  it  sublimes,  as  a 
white  crystalline  solid,  without  melting,  when  thus  prepared  it  is 
vaporized  and  condenses  in  a  cool  part  of  the  tube.  It  fumes 
when  exposed  to  moist  air  on  account  of  the  hydrogen  chloride 
produced  by  hydrolysis,  and  only  with  excess  of  hydrochloric  acid 
does  it  give  a  clear  solution  free  from  basic  salts. 

Aluminium  Hydroxide,  the  Aluminates  and  the  Oxide,  — 

When  an  alkali  is  added  to  a  solution  of  a  salt  of  aluminium,  the 
hydroxide  A1(OH)3  is  precipitated  in  gelatinous  form.  It  loses  water 
gradually  when  dried,  forming  no  intermediate  hydroxides,  until 
A12O3  remains.  Natural  forms  of  this  substance  are  hydrargyllite 
A1(OH)3  (=  A12O3,3H20),  bauxite  A120  (OH)  4  (=  A12O3,2H20),  which 
always  contains  ferric  oxide,  and  diaspore  A1O.OH  (=  A1203,H2O). 
Commercially,  the  hydroxide  is  made  by  heating  bauxite  with 
sodium  carbonate,  and  extracting  the  sodium  aluminate  with  water: 

A120(OH)4  +  Na2C03  ->  2NaA102  +  C02  +  2H20. 

The  hydroxide  is  then  precipitated  by  passing  carbon  dioxide 
through  the  solution: 

2NaAlO2  +  CO2  +  3H2O  ->  Na^CO,  +  2A1(OH)3. 


Aluminium  hydroxide  interacts  both  with  acids  and  with  bases, 
and  is,  therefore,  like  zinc  hydroxide  (p.  432),  ionized  both  as  a  base 
and  as  an  acid.  It  interacts  only  slightly  with  ammonium  hydroxide, 
because  this  substance  is  too  feebly  basic,  but,  from  the  solution  in 
the  active  alkalies,  the  aluminates  Na3.AlO3,  Na.AlO2,  and  K.A1O2, 
can  be  obtained  in  solid  form.  The  aluminates  are  largely  hydro- 
lyzed by  water: 

NaA102  4-  2H20  ^±  NaOH  +  A1(OH)3. 

Sodium  aluminate  NaA102  is  used  as  a  mordant   (see  below),  on 
account  of  the  ease  with  which  the  solution  gives  up  aluminium 


446  COLLEGE   CHEMISTRY 

hydroxide  when  any  material  is  present  which  can  combine  with  the 
free  portion  of  the  hydroxide  and  so  displace  forward  the  above 
equilibrium. 

When  calcium  chloride  is  added  to  a  solution  of  sodium  aluminate, 
the  insoluble  calcium  metaluminate  is  deposited: 

2NaAlO2  +  CaCl2  ->  Ca(A102)2  +  2NaCl. 

The  relations  of  these  various  substances  are  shown  by  the  following 
formulae: 

,0-H  /0-Na  < 

Al-O-H       Al-O-Na      Alf  Alf  ;Ca. 

^0-H  ^0-Na  °-H  °-Na    .,x° 


,  A  number  of  insoluble  metaluminates  are  found  in  nature.  They 
contain  bivalent  metals  in  place  of  the  calcium  in  the  last-named 
compound.  Thuswe  have  spinelleMg(AlO2)2,  and  gahnite  Zn(AlO2)2. 
Aluminium  oxide  A1203  (alumina)  is  found  in  nature  in  pure  form 
as  corundum.  This  mineral  is  only  one  degree  less  hard  than  the 
diamond.  Emery  is  a  common  variety,  contaminated  with  ferric 
oxide,  and  is  widely  used  as  an  abrasive.  The  ruby  is  pure  alu- 
minkpn  oxide  tinted  by  a  trace  of  a  compound  of  chromium,  while 
the  sapphire  is  the  same  material  colored,  possibly,  with  aluminate 
of  cobalt.  It  is  said,  however,  that  the  same  tint  is  conferred  upon 
colorless  corundum  by  exposure  to  the  influence  of  salts  of  radium. 

Aluminium  Sulphate:  The  Alums.  —  Aluminium  sulphate 
A12(S04)3,  18H20  is  prepared  by  treating  either  the  hydroxide  or 
pure  clay  (kaolin)  with  sulphuric  acid.  In  the  latter  case  the 
insoluble  residue  of  silicic  acid  is  removed  by  filtration: 

H2Al2(Si04)2  +  3H2S04-»  A12(S04)3  +  2H2Si03  +  2KA 

The  solution  of  the  sulphate  is  acid  in  reaction.  This  compound  is 
used  as  a  mordant  (see  below)  under  the  name  of  "  concentrated 
alum."  It  is  employed  also  in  sizing  cheaper  grades  of  paper,  an 
operation  required  to  prevent  the  absorption  and  consequent 
spreading  of  the  ink.  For  writing-paper,  gelatine  solution  is 
employed.  In  making  printing-papers,  rosin  soap  (made  by  dissolv- 


ALUMINIUM  447 

ing  rosin  in  caustic  soda)  is  mixed  with  the  pulp,  and  aluminium 
sulphate  is  added.  The  rosin  and  aluminium  hydroxide  are  precipi- 
tated in  the  pulp,  perhaps  in  feeble  combination,  and  pressing  be- 
tween hot  rollers  afterwards  melts  the  former  and  gives  a  surface 
to  the  paper. 

When  sulphate  of  potassium  is  added  to  a  strong  solution  of  alu- 
minium sulphate,  octahedral  crystals  of  potash  alum  (see  below) 
are  deposited.  This  is  a  double  salt,  and  is  one  of  a  large  class 
known  as  the  alums.  The  alums  have  the  general  formula  M2IS04, 
Af2m(S04)3,24H2O,  and  may  be  made  as  above  by  using  a  sulphate 
of  a  univalent  metal  with  one  of  a  trivalent  metal.  Thus,  for  M l  we 
may  use  K,  NH4,  Rb,  Cs,  and  Tl1,  and  for  M"1,  Al,  Fem,  Crm,  Mnm, 
and  Tlm.  All  the  alums  crystallize  in  octahedra. 

Potassium-aluminium  sulphate  K2SO4,A12(SO4)3,24H2O,  ordinary 
alum,  is  made  from  aluminium  sulphate.  It  is  also  prepared  by 
heating  alunite,  a  basic  alum  found  near  Rome  and  in  Hungary,  and 
extracting  the  product  with  water.  The  alunite  KAl3(OH)e(SO4)2 
leaves  an  insoluble  residue  of  the  hydroxide,  mixed  with  ferric  oxide 
which  is  present  as  an  impurity: 

2KA13(OH)8(SO4)2->K2SO4,A12(S04)3  +  4A1(OH)3. 

The  hydrated  salt  melts  at  90°.  An  aqueous  solution  of  this  salt, 
or  of  sodium  phosphate  (p.  380),  is  used  for  fire-proofing  draperies. 
The  crystals  deposited  in  the  fabric  melt  easily,  and  the  fused 
material  protects  the  fibers  from  access  of  oxygen.  When  heated 
more  strongly  alum  loses  its  water  of  hydration,  together  with  some 
sulphur  trioxide,  and  leaves  a  slightly  basic,  anhydrous  salt  known 
as  burnt  alum.  A  solution  of  alum  dissolves  a  considerable  amount 
of  aluminium  hydroxide,  giving  neutral  alum,  K2SO4,A14(OH)6(S04)3, 
a  basic  salt  used  as  a  mordant.  The  substance  is  usually  prepared 
by  adding  sodium  carbonate  to  the  solution  of  alum  as  long  as  the 
aluminium  hydroxide,  formed  locally,  continues  to  redissolve. 

Aluminium  sulphide  A12S3  is  most  easily  obtained  by  mixing  pyrite 
with  aluminium  powder  and  igniting  with  magnesium  ribbon 

(p  444V 

3FeS2  +  4A1  ->  2A12S3  +  3Fe. 

It  forms  a  grayish-black  solid,  and  is  decomposed  by  water,  like 
magnesium  sulphide,  giving  the  hydroxide  and  hydrogen  sulphide. 


448  COLLEGE   CHEMISTRY 

Dyeing:  Mordanting.  —  The  problem  of  the  dyer  is  to  confer  the 
desired  color  upon  a  fabric  made,  usually,  of  cotton,  linen,  wool,  or 
silk,  and  to  do  this  in  such  a  way  that  the  dye  is  fast  to  (i.e.,  is  not 
removed  or  destroyed  by)  rubbing,  and  often,  also,  to  washing  with 
soap.  To  understand  the  means  by  which  this  is  achieved,  it  must 
be  noted  that  cotton  and  linen  consist  of  hollow  fibers  of  the  com- 
position of  cellulose  (C6HwO5)a;.  Wool  is  made  of  hollow  fibers,  also, 
and  silk  of  rods,  but  the  material  is  entirely  different.  It  contains 
17  per  cent  of  nitrogen  in  the  case  of  wool,  and  20  per  cent  in  the  case 
of  silk,  and  the  nitrogen  compounds  of  which  the  material  is  com- 
posed are  much  more  active  chemically  than  is  cellulose,  and  com- 
bine incomparably  more  easily  and  firmly  with  the  many  kinds  of 
organic  compounds  which  are  used  as  dyes.  Hence,  stains  on  wool 
and  silk  are  much  less  often  removable  by  washing  than  are  those  on 
cotton. 

We  have  space  to  mention  only  three  kinds  of  dyes: 
1.  Insoluble  colored  bodies  which  are  formed  by  precipitation 
within  the  fibers  and  may  be  applied  to  any  fabric,  for  their  retention 
is  due  to  mechanical  and  not  to  chemical  causes.  If  cotton  is  boiled 
in  a  solution  of  lead  acetate  (or,  better  still,  sodium  plumbite,  q.v.), 
and  is  then  soaked  in  boiling  potassium  chromate  solution,  it  is  dyed 
a  brilliant  and  permanent  yellow.  Lead  chromate  is  the  colored 
body: 

Pb(C2H302)2  +  KA04^2KC2H302  +  PbCrOJ. 


In  indigo  dyeing  the  fabric  is  saturated  with  a  solution  of  indigo- 
white  in  caustic  soda,  and  is  then  exposed  to  the  air.  Indigo-blue  is 
formed  by  oxidation,  and,  being  insoluble,  is  precipitated  within  the 

fibers: 

2C18H12N202  +  02-»2C16H10N2021  +  2H20. 

2.  We  have  direct  or  substantive  dyes,  which  are  withdrawn  from 
a  solution  by  the  goods  which  are  being  dyed,  and  confer  upon  the 
latter  a  depth  of  color  depending  on  the  strength  of  the  solution  and 
the  affinity  of  the  material  for  the  dye.     These  dyes  are  fast  on  silk 
or  wool,  (for  example,  picric  acid)  but  only  a  small  minority  of  them 
are  taken  up  by  cotton  or  linen  in  such  a  way  that  they  cannot  be 
washed  out.     Congo  red,  C32H22N6S206Na2,  is  soluble  in  water,  and 
is  fast  both  on  cotton  and  on  wool. 

3.  The  last  class  comprises  the  mordant  or  adjective  dyes.     They 


ALUMINIUM  449 

work  on  the  principle  that  the  cloth  is  first  impregnated  with  a  sub- 
stance capable  of  attaching  itself  both  to  the  cloth  and,  subsequently, 
to  the  dye  also,  and  is  then  immersed  in  the  dye  itself.  Substances 
of  this  kind  are  tannic  acid  (for  basic  dyes)  and  colloidal  hydroxides 
(for  acid  dyes)  like  those  of  aluminium,  tin,  iron,  and  chromium. 
They  are  called  mordants  (Lat.  mordere,  to  bite).  When  aluminium 
hydroxide  is  to  be  used,  the  cloth  is  first  treated  with  a  hot  solution 
of  neutral  alum,  aluminium  sulphate,  or  sodium  aluminate,  and 
thereby  acquires,  either  by  adsorption  or  feeble  combination,  a 
certain  amount  of  the  hydroxide  (cf.  p.  445).  The  fabric  is  then 
boiled  in  water  with  the  dye.  If,  for  example,  alizarine  (madder)  is 
used,  the  cloth  is  dyed  Turkey  red.  Alizarine  is  an  orange-yellow, 
very  slightly  soluble  acid  of  the  composition  C14H8O4.  Since  the 
color  is  that  of  the  compound  of  the  dye  with  the  mordant,  different 
mordants  give  different  colors,  or  shades  of  color,  with  the  same  dye. 

Kaolin  and  Clay :  Earthenware  and  Porcelain.  —  By  the 

action  of  water  and  carbon  dioxide  upftn  granite  and  other  rocks 
containing  feldspar  KAlSi3O8,  the  potash  is  slowly  removed,  and 
the  compound  changed  largely  into  a  hydrated  orthosilicate  H2A12 
(SiO4)2,H20.  When  pure,  it  forms  kaolin  or  china  clay,  a  white, 
crumbly  material.  When  washed  away  and  redeposited,  it  usually 
acquires  compounds  of  iron,  and  the  carbonates  of  calcium  and 
magnesium,  becoming  common  clay.  Ocher,  umber,  and  sienna  are 
clays  colored  with  oxides  of  iron  and  manganese.  Fuller's  earth  is  a 
purer  variety. 

On  account  of  its  plasticity  when  moist,  and  its  tendency  to  become 
hard,  but  not  to  melt,  when  heated  strongly,  clay  is  used  in  making 
bricks,  pottery,  and  porcelain.  The  presence  of  calcium  and  magne- 
sium carbonates  makes  the  clay  more  fusible,  that  of  silica  less  so. 
Iron  compounds  cause  it  to  turn  red  during  firing.  For  earthenware, 
glazing  must  be  applied  to  make  the  vessels  water-tight.  This  is 
often  done  by  throwing  salt  into  the  kiln.  The  hot  steam  hydrolyzes 
the  salt  to  sodium  hydroxide  and  hydrochloric  acid,  and  the  former 
combines  with  the  clay,  giving  a  fusible  silicate  which  fills  the  pores 
of  the  surface.  For  porcelain,  very  pure  clay,  free  from  iron,  is 
employed,  and  it  is  mixed  with  feldspar  and  quartz.  The  feldspar 
melts  and  fills  the  pores  so  that  a  continuous,  semi-transparent 
material  results. 


450  COLLEGE   CHEMISTRY 

Analytical  Reactions   of  Aluminium    Compounds.  —  The 

alkalies,  and  alkaline  solutions  like  that  of  ammonium  sulphide, 
precipitate  the  white  hydroxide.  The  product  is  soluble  in  excess 
of  the  active  alkalies.  Soluble  carbonates  also  throw  down  the 
hydroxide.  Aluminium  compounds,  when  heated  strongly  in  the 
flame  with  cobalt  salts,  give  a  blue  aluminate  of  cobalt  Co(A102)2. 

Exercises.  —  1.   What  are  the  differences  between    zinc    and 
aluminium,  and  their  corresponding  compounds? 

2.  Construct  equations  showing,  (a)  the  hydrolysis  of  aluminium 
sulphate  (p.  446),  (b)  the  interaction  of  aluminium  sulphate  and 
cobalt  nitrate  in  the  Bunsen  flame. 

3.  Formulate  the  ionization  of  aluminium  hydroxide  (pp.  432, 445). 

4.  Why  does  zinc  hydroxide,  in  spite  of  its  feebleness  as  a  base, 
dissolve  in  ammonium  hydroxide,  while  aluminium  hydroxide  does 
not? 


CHAPTER  XXXIX 
GERMANIUM,    TIN,    LEAD 

THE  metallic  elements  of  the  fifth  column  of  the  periodic  table  are 
germanium  (Ge,  at.  wt.  72.5),  tin  (Sn,  at.  wt.  119),  and  lead  (Pb,  at. 
wt.  206.9).  These  are  on  the  right  side,  while  titanium  (Ti,  at.  wt. 
48.1),  zirconium  (Zr,  at.  wt.  90.6),  cerium  (Ce,  at.  wt.  140.25),  and 
thorium  (Th,  at.  wt.  232.5)  occupy  the  left  side. 

The  Chemical  Relations  of  the  Family.  —  All  of  these  ele- 
ments show  a  maximum  valence  of  four.  Germanium,  tin,  and  lead 

are  also  bivalent.     In  this  respect  they  resemble  carbon  and  differ 

^— «^_^  * 

from  silicon,  which  is  more  closely  allied  to  the  elements  on  the,  left 
side  of  the  column.  The  oxides  and  hydroxides  in  which  these  three 
elements  are  bivalent  become  more  basic,  and  the  elements  them- 
selves more  metallic  in  chemical  relations,  with  increase  in  atomic 
weight.  In  this  they  resemble  the  potassium,  calcium,  and  gallium 
families.  Curiously  enough,  the  same  three  hydroxides  are  also 
acidic.  They  are  more  strongly  acidic  than  is  zinc  hydroxide,  for 
the  salts  they  form  by  interaction  with  bases  are  less  hydrolyzed 
than  are  the  zincates.  This  acidic  character  likewise  increases  in  the 
order  in  which  the  elements  are  named  above. 

GERMANIUM. 

Germanium  (p.  277)  forms  two  oxides  GeO  and  GeO2  corresponding 
to  those  of  carbon  and  of  tin.  Germanious  oxide  isjiojtvery  definitely 
basic  or  acidic  T  and  the  sulphide  is  the  only  other~well-defined  com- 
pound of  this  set.  Germanic  oxide  and  hydroxide  are  acidic  entirely. 
The  resemblance  to  carbon  is  shown  in  the  formation  of  an  unstable 
compound  with  hydrogen,  of  germanium  chloroform  GeHCl3  and  of  a 
volatile  chloride  GeCl4  (b.-p.  87°). 

TIN. 

The  Chemical  Relations  of  the  Element.  —  Tin  is  both  biva- 
t  and  quadrivalent..  Each  of  the  oxides  and  hydroxides  SnO  and 
Sn(OH)2,  Sn62  and  SnO(OH)2  (or  Sn(OH)4),  is  both  basic  and  acidic, 

451 


452  COLLEGE   CHEMISTRY 

so  that  there  are  really  four  series  of  compounds.  Still,  stannous 
hydroxide  is  mainly  a  base,  of  a  feeble  sort,  while  stannic  hydroxide 
is  mainly  an  acid.  Thus  we  have  stannous  chloride,  sulphate,  and 
nitrate,  which  are  stable,  although  they  are  all  more  or  less  hydro- 
lyzed  by  water,  and  sodium  stannite  Na2.SnO2  which  is  unstable. 
On  the  other  hand,  stannic  nitrate,  sulphate,  and  chloride  are  com- 
pletely hydrolyzed  by  water,  while  sodium  stannate  Na2SnO3  is 
comparatively  stable.  The  dioxide  SnO2  is  an  infusible  solid, 
resembling  silicon  dioxide.  Tin  has  a  tendency  to  give  complex 
acids  and  salts,  like  H2SnCl6,  (NH4)2.SnCl6,  but  these  are  ionized 
also  to  a  small  extent  after  the  manner  of  double  salts,  giving  ions  of 
Sn"".  Tin  forms  no  salts  with  weak  acids,  like  carbonic  acid. 

Occurrence  and  Extraction.  —  The  chief  ore  of  tin  is  jin-stone, 
or  cassiteritp  ftnO»T  which  forms  square-prismatic  crystals  whose 
darTT color  is  due  to  the  presence  of  iron  compounds.  It  occurs  in 
Cornwall  and  the  East  Lidies.  The  ore  is  roughly  pulverized  and 
washed,  to  remove  granite  or  slate  with  which  it  is  mixed,  and  is 
then  roasted,  to  oxidize  the  sulphides  of  iron  and  copper,  and  drive 
off  the  arsenic  which  it  contains.  After  renewed  washing  to  elimi- 
nate sulphate  of  copper  and  oxide  of  iron,  it  is  reduced  with  coal  in  a 
reverberatory  furnace.  The  tin  is  afterwards  remelted  at  a  gentle 
heat,  and  the  pure  metal  flows  away  from  compounds  of  iron  and 
arsenic.  In  1900  the  production  was  4100  tons  and  63,700  tons  in 
England  and  in  the  East  Indies,  respectively.  These  quantities 
together  constitute  83  per  cent  of  the  world's  total  output. 

Physical  and  Chemical  Properties.  —  Tin  is  a  silver-white, 
crystalline  metal  of  low  tenacity  but  gj^t  malleability  (tinfoil).  Its 
specific  gravity  is  7.3,  and  its  meltingJPoint  about  233°. 

Tin-plate  is  made  by  dipping  carefully  cleaned  sheets  of  mild  steel 
into  molten  tin.  Vessels  of  copper  are  also  coated,  internally,  with 
tin,  to  prevent  the  formation  of  the  b^sic  carbonate  (p.  411).  For 
this  purpose  they  are  cleaned  with  ammonium  chloride,  sprinkled 
with  rosin  (to  reduce  the  oxide),  and  heated  to  230°.  Molten  tin  is 
then  spread  on  the  surface  with  a  piece  of  tow.  Alloys  of  tin,  such 
as  bronze  (p.  411),  soft  solder  (50  per  cent  lead),  pewter  (25  per  cent 
lead),  and  britannia  metal  (10  per  cent  antimony  and  some  copper), 
are  much  used  in  the  arts. 


TIN  453 

Tin,  although  it  displaces  hydrogen  from  dilute  acids,  is  not  tar- 
nished by  moist  air.  With  warm  hydrochloric  acid  it  gives  stannous 
chloride  SnCl2  and  hydrogen.  Hot,  concentrated  sulphuric  acid 
forms  stannous  sulphate  SnSO4  and  sulphur  dioxide  (cf.  p.  257). 
Nitric  acid,  when  cold  and  dilute,  interacts  with  it,  giving  stannous 
nitrate  Sn(NO3)2;  and  a  portion  of  the  nitric  acid  is  reduced  to 
ammonia  (cf.  p.  297). 

With  concentrated  nitric  acid,  stannic  nitrate  is  formed,  but  most 
of  this  salt  is  hydrolyzed  by  the  water  at  the  high  temperature  of  the 
action  (cf.  p.  438),  and  metastannic  acid  (H2SnO3)5  (a-stannic  acid) 
remains.  The  final  result  is  shown  by  the  equation  (simplified): 

Sn  +  4HN03  -»  H2SnO3  +  4NO2  +  H2O. 

Tin  also  displaces  hydrogen  from  caustic  alkalies,  giving  a  metastan- 
nate,  such  as  sodium  metastannate  Na2SnO3. 

Chlorides  of  Tin.  —  Stannous  chloride  SnCl2,  2H2O  is  made  by 
the  interaction pf  tin  and  hy drochloric ^acid .  When  the  crystals  are 
heated,  or  when  a  strong  aqueous  solution  is  diluted,  the  salt  is 
partially  hydrolyzed.  In  the  latter  case  the  basic  chloride  Sn (OH)C1 
is  deposited.  By  presence  of  excess  of  hydrochloric  acid,  the 
hydrolysis  is  prevented.  The  solution  is  used  as  a  mordant  (p.  449). 

Stannous  chloride  tends  to  pass  into  stannic  chloride  SnCl4,  and  is 
therefore  an  active  reducing  agent.  Thus,  it  reduces  the  chlorides  of 
mercury  (p.  437)  and  of  the  noble  metals,  liberating  the  free  metals. 
The  action  is  of  the  form  Hg"  +  Sn"  — >  Hg  +  Sn"**.  It  also  reduces 
free  oxygen,  or,  what  is  the  same  thing,  is  oxidized  by  the  air.  In 
this  case,  stannic  chloride  is  formed  in  the  acid  solution  and  the 
liquid  remains  clear;  in  the  neutral  solution  a  precipitate  of  the  basic 
chloride  is  formed  as  well:  ||j 

6SnCl2  4-  2H,O  +  O2  ->  4Sn(OH)Cl  +  2SnCl4. 

Powdered  tin,  if  placed  with  tjie  acid  solution,  will  undo  the  effects  of 
this  action  by  reducing  the  stannic  salt  to  the  stannous  condition. 

When  chlorine  acts  upon  tin,  or  upon  stannous  chloride  (either 
solid  or  dissolved),  stannic  chloride  SnCl4  is  formed.  The  compound 
is  a  colorless  liquid  (b.-p.  114°)  which  fumes  very  strongly  in  moist 
air,  giving  hydrochloric  acid  and  stannic  acid.  It  is  almost  com- 
pletely hydrolyzed  by  water.  The  stannic  acid  which  is  formed  is 


454  COLLEGE   CHEMISTRY 

not  precipitated,  however,  but  remains  dissolved  in  colloidal  (p.  96) 
SnCl4  +  4H2O^±  4HC1  +  Sn(OH)4. 


The  chloride,  with  small  amounts  of  water,  gives  hydrates,  of  which 
SnCl4,5H2O,  "  oxymuriate  of  tin,"  is  used  as  a  mordant.  Double 
(or  perhaps  complex)  salts,  such  as  ammonium-stannic  chloride  or 
"  pink-salt  "  (NH4)2SnCl6,  (used  as  a  mordant  on  cotton),  are 
readily  formed. 

Stannic  bromide  SnBr4  (b.-p.  201°)  resembles  stannic  chloride. 

OL-Stannic  Acid  and  its  Salts.  —  When  a  solution  of  stannic 
chloride  is  treated  with  ammonium  hydroxide,  a  white,  gelatinous 
precipitate  of  o-stannic  acid  is  formed: 

SnCl4  +  4NH4OH  -»  4NH4C1  +  H2SnO3  +  H2O. 

The  precipitate  loses  water  gradually  until  the  dioxide  remains,  and 
neither  Sn(OH)4  nor  SnO(OH)2  is  obtainable  as  a  definite  compound. 
When  stannic  oxide  is  fused  with  caustic  soda,  sodium  metastannate, 
or  a-stannate  Na-jSnOg^H-jO,  is  formed: 

SnO2  +  2NaOH  -»  Na.SnOg  +  H2O. 

This  compound  is  used  as  a  mordant  under  the  name  of  "  preparing 
salt."  When  its  solution  is  acidified,  the  above  mentioned  o-stannic 
acid  is  formed  by  double  decomposition.  This  ct-stannic  acid  inter- 
acts readily  with  acids  and  alkalies,  and  the  chloride  obtained  from 
it  is  identical  with  stannic  chloride  described  above. 

The  o-stannates  of  the  metals,  aside  from  those  of  potassium  and 
sodium,  like  the  silicates  and  carbonates  which  they  much  resemble, 
are  all  insoluble  in  water,  and  may  be  made  by  double  decomposition. 

(3-Stannic  Acid,  or  Metastannic  Acid.  —  The  product  of  the 
action  of  nitric  acid  upon  tin  (p.  453)  is  a  hydrated  stannic  oxide  like 
the  foregoing  substance,  but  is  not  identical  with  it.  It  is  not  easily 
soluble  in  alkalies.  By  boiling  it  with  caustic  soda,  however,  and 
then  extracting  with  pure  water,  a  soluble  sodiuni  /3-stannate  Na2Sn5Ou, 
is  obtained.  /8-stannic  acid  is  also  very  slowly  attacked  by  acids, 
and  the  chloride  secured  from  it  is  not  identical  with  the  ordinary 
chloride.  For  these  reasons  it  is  supposed  to  be  a  hydrate  of  a 
polymer  of  stannic  oxide  (Sn02)5,zH2O.  When  fused  with  caustic 
soda,  it  gives  the  same  a-stannate  as  does  the  dioxide  itself. 


TIN  455 

The  Oxides  of  Tin.  —  When  stannous  oxalate  is  heated  in 
absence  of  air,  stannous  oxide  SnO  remains:  SnC2O4— >SnO  +  CO2  +  CO.' 
It  is  a  black  powder  which  burns  in  the  air,  giving  the  dioxide.  The 
corresponding  hydroxide  Sn2O(OH)2  is  formed  by  adding  sodium 
carbonate  to  stannous  chloride  solution.  It  is  a  white  powder,  easily 
dehydrated,  and  interacts  with  alkalies  to  give  soluble  stannites, 
such  as  Na2SnO2.  With  acids,  the  hydroxide  gives  stannous  salts. 

Stannic  oxide  Sn02  is  found  in  nature  (p.  452),  and  may  be  made 
in  pure  form  by  igniting  /8-stannic  acid.  When  heated,  it  becomes 
yellow,  but  recovers  its  whiteness  when  cooled  (cf.  Zinc  oxide,  p.  432). 
Prepared  at  a  low  temperature,  it  interacts  easily  with  acids,  but 

after  strong  ignition,  is  affected  by  them  very  slowly. 

( 

The  Sulphides  •/  &n.  —  Stannous  sulphide  SnS  is  obtained  as 
a  dark-brown  precipitate  when  hydrogen  sulphide  is  led  into  a 
solution  of  a  stannous  salt. 

Stannic  sulphide  SnS2  is  formed  likewise  by  precipitation,  and  is 
yellow  in  color.  Stannic  sulphide  loses  sulphur  when  strongly 
heated,  and  leaves  stannous  sulphide.  It  is  not  much  affected  by 
dilute  acids,  but  interacts  with  solutions  of  ammonium  sulphide 
(or  sodium  sulphide),  giving  a  soluble  complex  sulphide,  namely, 
ammonium  sulphostannate : 

SnS2  +  (NH4)2S  ->  (NH4)2.SnS3. 

The  corresponding  sodium  sulphostannate  is  easily  crystallized  in  the 
form  Na2SnS3,2H2O.  Stannous  sulphide  is  not  affected  by  soluble 
sulphides,  but  polysulphides,  such  as  yellow  ammonium  sulphide, 
give  with  it  the  above  mentioned  sulphostannates: 

SnS  +  (NH4)2S  +  S  ->  (NH4)2.SnS3. 

With  acids  the  sulphostannates  undergo  double  decomposition,  but 
the  free  acid  H2.SnS3  thus  produced  is  unstable  and  breaks  up, 
giving  off  hydrogen  sulphide,  and  depositing  stannic  sulphide. 

Analytical  Reactions  of  Salts  of  Tin.  —  The  two  ionic 
forms  of -tin,  Sn",  and  Sn""*,  are  both  colorless.  Their  behavior  is 
different.  They  give  a  brown  and  a  yellow  sulphide,  respectively, 
with  hydrogen  sulphide.  These  sulphides  dissolve  in  yellow  ammo- 
nium sulphide  (above).  The  reducing  power  of  stannous-ion  Sn" 
is  very  characteristic  (p.  453).  The  oxides  are  reduced  by  charcoal 
in  the  reducing  part  of  the  Bunsen  flame  and  the  metal  is  liberated. 


456  COLLEGE   CHEMISTR 

LEAD. 

The  Chemical  Relations  of  the  Element.  —  Lead  is  both 
bivalent  and  quadrivalent.  The  oxides  PbO  and  PbO2,  and  the 
corresponding  hydrated  oxides,  are  all  both  basic  and  acidic.  Lead 
monoxide  is  a  fairly  active  base,  comparable  with  cupric  oxide,  but 
lead  dioxide  is  a  feeble  one.  Both  are  feebly  acidic.  The  salts  of 
bivalent  lead,  like  Pb(NO3)2,  commonly  called  the  plumbic  salts,  are 
somewhat  hydrolyzed  by  water*  but  less  so  than  are  those  of  tin. 
The  tetrachloride  and  other  salts  of  quadrivalent  lead  are  completely 
hydrolyzed.  The  plumbites  Na<>.PbO2  and  plumbates  Na2.PbO3  are 
hydrolyzed  to  a  considerable  extent.  All  the  compounds  in  which 
lead  is  quadrivalent  give  up  half  of  the  negative  radical  readily,  and 
are  reduced  to  the  "  plumbic  "  condition.  The  metal  displaces 
hydrogen  with  difficulty,  and  is  easily  displaced  by  zinc.  Lead  com- 
pounds are  all  poisonous,  and  the  effects  of  repeated,  very  minute 
doses  are  cumulative, —  resulting  in  "  lead  colic." 

Occurrence  and  Metallurgy.  —  Commercial  lead  is  almost  all 
obtained  from  galena  PbS,  which  crystallizes  in  cubes.  This  ore 
often  contains  considerable  amounts  of  silver  sulphide  Ag2S. 

The  sulphide  of  lead  is  first  roasted  until  a  sufficient  proportion 
of  it  has  been  converted  into  the  oxide  and  sulphate.  The  furnace- 
doors  are  then  closed,  and  the  temperature  raised  in  order  that  these 
products  may  interact  with  the  unchanged  part  of  the  sulphide: 

PbS  +  2PbO  ->  3Pb  +  S00, 
PbS  +  PbS04  -»  2Pb  +  2S02. 

Another  plan  consists  in  heating  galenite  with  scrap  iron  or  iron  ores 
and  coal:  PbS  +  Fe— »Pb  +  FeS.  The  molten  ferrous  sulphide 
rises  to  the  top  as  a  matte. 

Physical  and  Chemical  Properties.  —  Metallic  lead  is  gray 
in  color,  very  soft,  and  of  small  tensile  strength.  Its  specific  gravity 
is  11.4,  and  its  melting-point  326°.  While  warm,  it  is  formed  by 
hydraulic  pressure  into  pipes  which  are  used  in  plumbing  and  for 
covering  electric  cables.  On  account  of  its  very  slow  interaction 
with  most  substances,  sheet  lead  is  used  in  chemica^actories,  for 
example,  to  line  sulphuric-acid  chambers.  An  alloy^Pntaining  0.5 
per  cent  of  arsenic  is  used  in  making  small  shot  and  shrapnel  bullets. 


LEAD  457 

Type-metal  contains  20-25  per  cent  of  antimony  (q.v.).  In  both 
cases  greater  hardness  is  secured  by  the  addition  of  the  foreign  metal. 

Lead  oxidizes  very  superficially  in  the  air.  The  suboxide  Pb2O  is 
supposed  to  be  first  formed.  The  final  covering  is  a  basic  carbonate. 
Contact  with  hard  waters  confers  upon  lead  a  similar  coating  com- 
posed of  the  carbonate  and  the  sulphate.  These  deposits,  being 
insoluble,  inclose  the  metal  and  protect  the  water  from  contamina- 
tion with  lead  compounds.  Pure  rain-water,  however,  since  it  has 
no  hardness,  and  contains  oxygen  in  solution,  gives  the  hydroxide 
Pb(OH)2,  which  is  noticeably  soluble.  When  heated  in  the  air, 
lead  gives  the  monoxide  PbO  or  minium  Pb3O4,  the  latter  at 
lower  temperatures. 

The  metal  displaces  hydrogen  from  hydrochloric  acid  slowly.  It 
is  hardly  affected  by  concentrated  sulphuric  acid  (cf.  p.  262).  Nitric 
acid  attacks  it  readily,  giving  lead  nitrate  and  oxides  of  nitrogen 
(p.  298). 

Chlorides  and  Iodide.  —  Plumbic  chloride  PbCl2  is  precipitated 
when  a  soluble  chloride  is  added  to  a  solution  of  a  lead  salt.  It  is 
slightly  soluble  in  water  (1.5  :  100)  at  18°,  and  much  more  so  at  100°. 

Lead  tetrachloride  PbCl4  is  a  solid  at  —  15°,  and  loses  chlorine  at 
the  ordinary  temperature.  It  is  made  by  passing  chlorine  into 
plumbic  chloride  suspended  in  hydrochloric  acid.  The  solution 
appears  to  contain  H2PbCl6.  When  this  is  thrown  into  cold,  con- 
centrated sulphuric  acid,  an  oil,  PbCl4,  settles  to  the  bottom.  The 
oil  fumes  in  the  air,  and  closely  resembles  stannic  chloride  SnCl4. 
With  little  water,  it  slowly  deposits  PbCl2  and  gives  off  chlorine. 
With  much  water  it  is  quickly  hydrolyzed,  and  lead  dioxide  is 
thrown  dowft:  PbCl4  +  2H2O  ->  PbO2  +  4HC1. 

The  yellow  lead  iodide  PbI2  is  formed  by  precipitation.  It  crys- 
tallizes in  yellow  scales  from  solution  in  hot  water. 

Oxides  and  Hydroxides.  —  There  are  five  different  oxides  of 
lead,  Pb20,  PbO,  Pb3O4,  Pb2O3,  and  PbO2.  The  suboxide  Pb2O  is  a 
dark-gray  powder,  formed  by  gently  heating  the  oxalate.  Plumbic 
oxide,  or  lead  monoxide  PbO,  is  made  by  cupellation  (p.  418)  of 
lead,  and  the  solidified,  crystalline  mass  of  yellowish-red  color  is 
sold  as  "  litharge."  All  the  other  oxides  yield  this  one  when  they 
are  heated  above  600°  in  the  air.  Plumbic  oxide  takes  up  carbon 


458  COLLEGE  CHEMISTRY 

dioxide  from  the  air,  and  therefore  usually  contains  a  basic  carbonate. 
The  oxide  is  used  in  glass-making  and  for  preparing  salts  of  lead. 

Plumbic  hydroxide  Pb(OH)2  is  formed  by  precipitation.  It  gives  up 
water  in  stages,  the  successive  products  being  Pb(OH)2,  Pb2O(OH)2, 
Pb3O2(OH)2.  These  substances  are  equivalent  in  composition  to 
PbO,H2O,  2PbO,H2O,  and  3PbO,H2O  respectively.  The  hydroxide 
is  observably  soluble  in  water,  and  gives  a  solution  with  a  faintly 
alkaline  reaction.  With  acids  it  forms  salts  of  lead.  It  interacts 
also  with  potassium  and  sodium  hydroxides  to  form  the  soluble 
plumbites,  like  sodium  plumbite  Na2.PbO2. 

Minium,  or  red  lead,  Pb3O4,  gives  off  oxygen  when  heated: 

2Pb304  <=>  6PbO  +  02. 

On  account  of  unequal  heating  during  manufacture,  commercial  red 
lead  is  never  fully  oxidized,  and  always  contains  litharge.  Con- 
versely, commercial  litharge  usually  contains  a  little  minium. 

Minium,  when  heated  with  warm,  dilute  nitric  acid,  is  decomposed, 
and  leaves  lead  dioxide  as  an  insoluble  powder.  It  is  therefore 
regarded  as  lead  orthoplumbate  (see  below) : 

Pb2.PbO4  +  4HNO3^2Pb(NO3)2  +  H4Pb04. 

The  double  decomposition  as  a  salt  that  it  thus  undergoes  is  followed 
by  dehydration  of  the  plumbic  acid,  which  is  unstable  (H4PbO4  — > 
PbO2  +  2H20),  and  the  dioxide  remains.  Red  lead  is  used  in 
glass-making,  and,  when  mixed  with  oil,  gives  a  red  paint. 

Lead  dioxide  PbO2  may  be  obtained  as  described  above  in  the  form 
of  a  brown  powder.  It  is  usually  made  by  adding  bleaching  powder 
to  an  alkaline  solution  of  plumbic  hydroxide: 

Na^.PbO,  +  Ca(OC!)Cl  +  H2O  ->  2NaOH  +  CaCl2  +  Pb02  J'. 

In  this  action  we  may  regard  the  free  lead  hydroxide,  formed  by 
hydrolysis  of  the  plumbite,  as  being  oxidized  by  the  bleaching  pow- 
der. Lead  dioxide  is  an  active  oxidizing  agent.  It  interacts  with, 
and  sets  fire  to,  a  stream  of  hydrogen  sulphide,  and  it  liberates 
chlorine  from  hydrochloric  acid.  With  acids  it  gives  no  hydrogen 
peroxide,  and  is  not  a  peroxide  in  the  restricted  sense  of  the  term 
(p.  212).  Lead  dioxide  interacts  with  potassium  and  sodium  hydrox- 
ides, giving  soluble  plumbates.  The  potassium  salt  K2Pb03.3H20  is 
analogous  to  the  metastannate  K2SnO3,3H20  (p.  454).  A  mixture 


LEAD  459 

* 

of  calcium  carbonate  and  lead  monoxide  absorbs  oxygen  when  heated 
in  a  stream  of  air,  and  the  yellowish-red  calcium  orthoplumbate  is 
formed : 

4CaCO3  +  2PbO  +  O2  <=>  2Ca2PbO4  +  4CO2. 

The  action  is  reversible,  and  is  at  the  basis  of  Kassner's  method  of 
manufacturing  oxygen  from  the  air. 

Other  Salts  of  Lead.  —  Lead  nitrate  Pb(NO3)2  may  be  made 
by  treating  lead,  lead  monoxide,  or  lead  carbonate  with  nitric  acid. 
It  forms  white,  anhydrous  octahedra.  The  nitrate  and  acetate  (see 
below)  are  the  salts  of  lead  which,  because  of  their  solubility  (see 
Table),  are  most  commonly  used.  On  account  of  hydrolysis,  the 
solution  of  the  nitrate  is  acid  in  reaction. 

Lead  carbonate  PbCO3  is  found  in  nature.  It  may  be  formed  as  a 
precipitate  by  adding  a  soluble  bicarbonate  to  lead  nitrate  solution. 
With  normal  sodium  carbonate,  a  basic  carbonate  Pb3(OH)2(CO3)2 
is  deposited.  This  basic  salt  is  identical  with  white  lead,  which,  on 
account  of  its  superior  opacity,  has  better  covering  power  than  zinc- 
white  (p.  432)  or  permanent  white  (p.  405).  The  substance  is  man- 
ufactured in  various  ways,  all  of  which  involve  the  oxidation  of  the 
lead  by  the  air,  the  formation  of  a  basic  acetate  by  the  interaction  of 
vinegar  or  acetic  acid  with  the  oxide,  and  the  subsequent  decompo- 
sition of  the  salt  by  carbon  dioxide.  The  best  quality  is  obtained 
by  the  Dutch  method.  In  this,  gratings  of  cast  lead  are  placed 
above  a  shallow  layer  of  vinegar  in  small  pots.  These  pots  are 
buried  in  manure,  which  by  its  decomposition  furnishes  the  carbon 
dioxide  and  the  necessary  warmth.  The  gratings  are  gradually  con- 
verted into  a  white  mass  of  the  basic  carbonate.  The  vapor  of  acetic 
acid  arising  from  the  vinegar  may  be  regarded  as  a  catalytic  agent 
(cf.  p.  54),  since  it  is  used  over  and  over  again. 

Lead  acetate  Pb(C2H3O2)2,3H20  is  made  by  the  action  of  acetic  acid 
on  litharge.  It  is  easily  soluble  in  water,  and,  from  the  sweet  taste 
of  the  solution,  is  named  sugar  of  lead  (used  in  medicine).  The 
basic  salt  Pb(OH)(C2H3O2)  is  formed  by  boiling  a  solution  of  lead 
acetate  with  excess  of  litharge.  Unlike  most  basic  salts,  this 
basic  salt  is  soluble  in  water,  and  its  solution  has  a  faintly  alkaline 
reaction. 

Lead  sulphate  PbSO4  occurs  in  nature  as  anglesite.  Being  insol- 
uble in  water,  it  is  easily  obtained  by  precipitation. 


460  COLLEGE  CHEMISTRY 

Natural  lead  sulphide  PbS  (galena)  is  black,  and  its  crystals  have 
a  silvery  luster.  The  precipitated  salt  is  amorphous.  It  is  more 
easily  attacked  by  active  acids  than  is  mercuric  sulphide  (cf.  p.  438). 

Analytical  Reactions  of  Lead  Compounds.  —  Hydrogen 
sulphide  precipitates  the  black  sulphide,  even  when  dilute  acids  are 
present.  Sulphuric  acid  throws  down  the  sulphate.  Potassium 
hydroxide  gives  the  white  hydroxide,  which  dissolves  in  excess  to 
form  the  plumbite.  Potassium  chromate  or  dichromate  (q.v.)  gives 
a  yellow  precipitate  of  lead  chromate  PbCrO4,  which  h  used  as  a 
pigment  under  the  name  of  "  chrome-yellow." 

TITANIUM,  ZIRCONIUM,  CERIUM,  THORIUM. 

The  metals  on  the  left  side  of  the  fifth  column  of  the  periodic  table 
are  all  quadrivalent,  although  compounds  in  which  a  lower  valence 
appears  are  numerous  in  this  family.  The  first  two  are  feebly  base- 
forming  as  well  as  feebly  acid-forming;  the  last  two  are  base-forming 
exclusively. 

Titanium  occurs  in  rutile  TiTi04.  Derived  from  it  are  a  number  of 
titanates  of  the  form  K2TiO3.  Zirconium  is  found  in  zircon,  the 
orthosilicate  of  zirconium  ZrSiO4.  The  oxide  is  used  in  making  the 
incandescent  substance  in  some  forms  of  gas  lamps. 

Cerium  occurs  chiefly  in  cerite  [Ce,  La,  Nd,  Pd]  SiO4,H20  (cf.  p. 
443). 

Thorium  is  found  in  thorite  ThSiO4,  but  most  of  the  supply  comes 
from  monazite  sand.  The  nitrate  Th(NO3)4,6H2O  is  used  in  making 
Welsbach  incandescent  mantles  (cf.  Flame,  p.  338). 

The  Nernst  lamp  is  an  incandescent  electric  lighting  arrangement 
in  which  a  rod  of  the  oxides  of  several  of  the  rare  metals  takes  the 
place  of  the  common  carbon  filament.  The  peculiarity  of  this  lamp 
is  that  preheating  is  required  before  the  rod  attains  a  temperature  at 
which  it  will  conduct  the  current.  When  this  point  has  been  once 
reached,  the  resistance  enables  the  current  to  maintain  the  rod  at 
the  temperature  of  incandescence.  For  equal  consumptions  of 
electricity,  this  form  of  electric  lamp  gives  a  greater  yield  of  light 
than  does  the  ordinary,  carbon-filament,  incandescent  bulb. 

Exercises.  —  1 .  In  what  order  should  you  place  the  elements 
dealt  with  in  this  chapter,  beginning  with  the  least  metallic,  and  end- 
ing with  the  most  metallic  (p.  353)  ? 


TITANIUM,  ZIRCONIUM,  CERIUM,  THORIUM  461 

2.  Construct  equations  showing,  (a)  the  interaction  of  tin  and  con- 
centrated  sulphuric  acid,   (b)  of  water  and  stannous  chloride,  (c) 
of  oxygen  and  stannous  chloride  in  acid  solution,  (d)  the  decom- 
position of  lead  oxalate  (p.  457),  (e)  the  interaction  of  lead  mon- 
oxide and  acetic  acid,  (/)  and  of  lead  monoxide  and  lead  acetate. 

3.  To  which  class  of  ionic  actions  (pp.  243,  253,  412)  do  the  reduc- 
tions by  stannous  chloride  and  by  tin  (p.  453)  belong? 

4.  What  interactions  probably  occur  when  lead  dioxide  liberates 
chlorine  from  hydrochloric  acid? 

5.  How  should  you  set  about  preparing,   (a)  lead  oxalate  (in- 
soluble), (b)  lead  chlorate  (soluble)? 


CHAPTER  XL 

e  £-?•'" 
ARSENIC,    ANTIMONY,   BISMUTH 

THIS  family  is  very  closely  related  to  the  elements  phosphorus  and 
nitrogen  which  precede  it  in  the  same  column  of  the  periodic  table. 
In  reading  this  chapter,  therefore,  constant  reference  should  be  made 
to  the  chemistry  of  the  corresponding  compounds  of  phosphorus. 
For  a  general  comparison  of  the  elements  arsenic  (As,  at.  wt.  75), 
antimony  (Sb,  at.  wt.  120.2)  and  bismuth  (Bi,  at.  wt.  208)  with  each 
other  and  with  the  two  already  disposed  of,  see  p.  471.  It  is  sufficient 
here  to  say  that  arsenic  is  mainly  an  acid-forming  element,  and  is 
therefore  a  non-metal,  while  antimony  is  both  acid-forming  and 
base-forming,  and  bismuth  is  base-forming.  Each  of  the  three 
elements  gives  two  sets  of  compounds,  in  which  it  is  trivalent,  and 
quinquivalent,  respectively.  None  of  the  free  elements  displaces 
hydrogen  from  dilute  acids. 

ARSENIC  As. 

The  Chemical  Relations  of  the  Element.  —  Arsenic  forms  a 
compound  with  hydrogen  AsH3.  It  gives  several  halogen  derivatives 
of  the  type  AsX3  which  are  completely  hydrolyzed  by  water.  Its 
oxides  and  hydroxides  are  acidic. 

Sulphates,  nitrates,  carbonates,  and  other  salts  of  arsenic  are  not 
formed.  The  complex  sulphides  (p.  455)  are  important. 

Occurrence  and  Preparation. — Arsenic  is  found  free  in  nature. 
It  occurs  also  in  combination  with  many  metals,  particularly  in 
arsenical  pyrites  FeAsS.  Two  sulphides  of  arsenic,  orpiment  As2S3 
and  realgar  As2S2,  and  an  oxide  As2O3,  are  less  common. 

The  element  is  obtained  either  from  the  native  material  or  by 
heating  arsenical  pyrites:  FeAsS  — >  FeS  +  As.  During  the  roasting 
of  the  sulphur  ores  of  metals,  arsenic  trioxide  is  formed  by  the  oxida- 
tion of  the  arsenic  so  frequently  present,  and  collects  as  a  dust  in  the 
flues. 

462 


ARSENIC  463 

Physical  and  Chemical  Properties. — The  free  element  is  steel- 
gray  in  color,  metallic  in  appearance,  and  crystalline  in  form.  It  is 
easily  volatilized  at  180°,  and  acquires  a  vapor  pressure  of  760  mm. 
long  before  the  melting-point  (480°,  under  high  pressure)  is  reached. 
The  density  of  the  vapor  measured  at  644°  gives  308.4  as  the  weight 
of  the  G.M.V.  (22.4  liters  at  0°  and  760  mm.).  The  weight  of  arsenic 
combining  with  one  chemical  unit  weight  (35.45  g.)  of  chlorine,  is 
25  g.  Three  times  this  amount,  or  75  g.,  is  the  smallest  weight  found 
in  the  G.M.V.  of  any  volatile  compound  of  arsenic,  and  is  therefore 
accepted  as  the  atomic  weight  (p.  134).  Since  308.4  is  equal  approxi- 
mately to  4  X  75  (  =  300),  the  formula  of  the  vapor  of  the  simple 
substance  at  644°  is  As4.  At  1700°  the  formula  is  As2  (cf.  p.  146). 

The  free  element  burns  in  the  air,  producing  clouds  of  the  solid 
trioxide  As2O3.  It  unites  directly  with  the  halogens,  with  sulphur, 
and  with  many  of  the  metals.  When  boiled  with  nitric  acid,  chlorine 
water,  and  other  powerful  oxidizing  agents  (p.  194),  it  is  oxidized  in 
the  same  way  as  is  phosphorus,  and  yields  arsenic  acid  H3AsO4. 


Arsine  AsH^.  —  This  substance  corresponds  in  composition  to 
ammonia  and  phosphine,  and  some  of  the  ways  in  which  it  may  be 
formed  are  analogous  to  those  used  in  the  case  of  these  substances. 
Thus,  when  arsenic  and  zinc  are  melted  together  in  the  proportions 
to  form  zinc  arsenide  Zn3As2,  and  the  product  is  treated  with  dilute 
hydrochloric  acid,  the  result  is  similar  to  the  action  of  water  or  dilute 
acids  upon  calcium  phosphide,  and  arsine  is  evolved  as  a  gas: 

6HC1  -»  2AsH3  +  3ZnCl2. 


Arsine  (arsenuretted  hydrogen)  is  formed  also  by  the  action  of  nas- 
cent hydrogen  (cf.  p.  302)  upon  soluble  compounds  of  arsenic.  When 
a  solution  of  arsenious  chloride  AsCl3  or  arsenic  acid  is  added  to 
zinc  and  hydrochloric  acid  in  a  generating  flask,  arsine  is  formed: 

AsCl3  +  3H2  -*  AsH3  +  3HC1. 

Pure  arsine  may  be  secured  by  leading  this  mixture  with  hydrogen 
through  a  U-tube  immersed  in  liquid  air.  The  arsine  (b.-p.  —  40°) 
condenses  as  a  colorless  liquid. 

Arsine  burns  with  a  bluish  flame,  producing  water  and  clouds  of 
arsenic  trioxide:  2AsH3  +  3O2  —  >  3H2O  +  As2O3.  The  combustion 
of  hydrogen  containing  arsine,  generated  as  just  described,  gives  the 


464  COLLEGE  CHEMISTRY 

same  substances.  Since  arsine,  when  heated,  is  readily  dissociated 
into  its  constituents  (cf.  p.  251),  the  vapor  of  free  arsenic  is  present 
in  the  interior  of  the  hydrogen  flame.  This  arsenic  may  be  con- 
densed in  the  form  of  a  metallic-looking,  brownish  stain  by  inter- 
position of  a  cold  vessel  of  white  porcelain.  Even  when  only  a  trace 
of  the  compound  of  arsenic  has  been  added  to  the  materials  in  the 
generator,  the  stain  which  is  produced  is  very  conspicuous.  This 
behavior  thus  furnishes  us  with  the  basis  of  an  exceedingly  delicate 
test  —  Marsh's  test  —  for  the  presence  of  arsenic  in  any  soluble 
form  of  combination.  The  compounds  of  antimony  alone  show  a 
similar  phenomenon  (see  Stibine). 

Arsine  .is  exceedingly  poisonous,  the  breathing  of  small  amounts 
producing  fatal  effects.  It  differs  from  ammonia  more  markedly 
than  does  phosphine,  for  it  is  not  only  without  action  on  water  or 
acids,  but  does  not  unite  directly  even  with  the  halides  of  hydrogen. 

Halides  of  Arsenic*  —  The  halides  include  a  liquid  trifluoride 
AsF3,  a  liquid  trichloride,  a  solid  tribromide  AsBr3,  and  a  solid  tri- 
iodide  AsI3. 

The  trichloride  AsCl3,  which  is  prepared  by  passing  chlorine  gas 
into  a  vessel  containing  arsenic,  is  easily  formed  as  the  result  of  a 
vigorous  action.  It  is  a  colorless  liquid  (b.-p.  130°).  When  mixed 
with  water  it  is  at  once  converted  into  the  white,  almost  insoluble 
trioxide.  The  action  is  presumably  similar  to  that  of  water  upon 
the  corresponding  compound  of  phosphorus  (p.  163),  but  the  arseni- 
ous  acid  for  the  most  part  loses  water  and  forms  the  insoluble 
anhydride: 

AsCl3  +  3H2O  <F±  As(OH)3  +  3HC1, 
2  As  (OH)  3  ^±  As2O3l  +  3H20. 

This  action,  however,  differs  markedly  from  the  other  in  that  it  is 
reversible,  and  arsenic  trioxide  interacts  with  aqueous  hydrochloric 
acid,  giving  a  solution  of  arsenious  chloride. 

Oxides  of  Arsenic.  —  Arsenic  trioxide  As2O3  is  produced  by 
burning  arsenic  in  the  air  and  during  the  roasting  of  arsenical  ores 
(p.  462),  and  is  known  as'  "  white  arsenic  "  or  simply  "  arsenic." 
It  is  purified  for  commercial  purposes  by  subliming  the  flue-dust  in 
cylindrical  pots.  The  pure  trioxide  is  deposited  in  a  glassy  form 
in  the  upper  part  of  the  vessel. 


ARSENIC  465 

When  treated  with  water,  the  trioxide  goes  into  solution  to  a 
very  slight  extent  (0.3: 100),  forming  arsenious  acid,  by  reversal  of 
the  second  of  the  actions  given  above.  In  boiling  water  the  solu- 
bility is  greater  (11.5  : 100).  When  heated  in  a  tube  with  carbon,  this 
oxide  is  reduced,  and  the  free  element,  being  volatile,  is  deposited 
upon  the  cold  part  of  the  tube  just  above  the  flame.  The  trioxide 
is  an  active  poison,  since  it  gradually  passes  into  solution,  forming 
arsenious  acid. 

The  pentoxide  As2O5  is  a  white  crystalline  substance,  formed  by 
heating  arsenic  acid:  2H3AsO4  — >  As2O5  +  3H2O.  When  raised  to 
a  higher  temperature,  it  loses  a  part  of  its  oxygen,  leaving  the 
trioxide.  In  consequence  of  this  instability,  it  cannot  be  formed  by 
direct  union  of  oxygen  with  the  trioxide,  after  the  manner  of  phos- 
phorus pentoxide. 

Acids  of  Arsenic.  —  When  elementary  arsenic  or  arsenious 
oxide  is  treated  with  concentrated  nitric  acid,  or  with  chlorine  and 
water,  orthoarsenic  acid  H3AsO4  is  produced.  The  substance,  a 
deliquescent  white  solid,  possesses  a  composition  similar  to  that  of 
orthophosphoric  acid  and,  like  the  latter  (p.  309),  when  heated  loses 
water  in  progressive  stages,  furnishing  intermediate  acids  —  pyro- 
arsenic  acid  and  metarseidc  acid  —  and  finally  the  pentoxide. 
relationship  of  the  substances  is  shown  by  the  formulae: 

H3As04  — >  H4As207  ->  HAs03  ->  As206. 


-  pyro- 
.    The 


These  acids  differ  from  the  corresponding  compounds  of  phosphorus 
in  that  upon  solution  in  water  they  immediately  pass  back  into  the 
ortho-acid.  With  metaphosphoric  acid,  also,  the  final  elimination 
of  all  the  water  by  simple  heating  is  impossible.  The  chocolate- 
brown  silver  orthoarsenate  Ag3AsO4  and  the  white  MgNH4As04,  like 
the  corresponding  phosphates,  are  insoluble  in  water. 

Arsenious  acid  H3AsO3,  like  sulphurous  and  carbonic  acids,  loses 
water,  and  yields  the  anhydride  (arsenic  trioxide)  when  the  attempt 
is  made  to  obtain  it  from  the  aqueous  solution.  The  potassium  and 
sodium  arsenites,  K3AsO3  and  Na3AsO3,  are  made  by  treating  arsenic 
trioxide  with  caustic  alkalies,  and  are  much  hydrolyzed  by  water. 
The  arsenites  of  the  heavy  metals  are  insoluble,  and  can  be  made  by 
precipitation.  Paris  green  (p.  415)  is  an  arsenite  of  copper.  In 
cases  of  poisoning  by  white  arsenic,  freshly  precipitated  ferric 


466  COLLEGE   CHEMISTRY 

hydroxide  (or  the  same  compound  in  colloidal  solution)  or  magne- 
sium hydroxide  is  administered,  since  by  interaction  with  the 
arsenious  acid  they  form  insoluble  arsenates  and  arsenites. 

Sulphides  of  Arsenic,  —  Arsenic  pentasulphide  As2S5  is  obtained 
as  a  yellow  powder  by  decomposition  of  the  sulpharsenates  (see 
below),  and  by  leading  hydrogen  sulphide  into  a  solution  of  arsenic 
acid  in  concentrated  hydrochloric  acid. 

Arsenious  sulphide  As2S3  occurs  in  nature  as  orpiment,  and  was 
formerly  used  as  a  yellow  pigment  (auripigmentum).  The  word 
arsenic  is  derived  from  the  Greek  name  for  this  mineral  (d/ao-eviKoV). 
It  is  obtained  as  a  citron-yellow  precipitate  when  hydrogen  sulphide 
is  led  into  an  aqueous  solution  of  arsenious  chloride. 

Realgar  As2S2  is  a  natural  sulphide  of  orange-red  color,  and  is  also 
manufactured  by  subliming  a  mixture  of  arsenical  pyrites  and  pyrite : 

2FeAsS  +  2FeS2-»4FeS  +  As2S2t 

It  burns  in  oxygen,  forming  arsenious  oxide  and  sulphur  dioxide,  and 
is  mixed  with  potassium  nitrate  and  sulphur  to  make  "  Bengal 
lights." 

Sulpharsenites  and  Sulpharsenates.  —  The  sulphides  of 
arsenic  interact  with  solutions  of  alkali  sulphides  after  the  manner 
of  the  sulphides  of  tin  (p.  455),  giving  soluble,  complex  sulphides. 
Arsenious  sulphide  with  colorless  ammonium  sulphide  gives  ammo- 
nium sulpharsenite,  and  with  the  yellow  sulphide  gives  ammonium 
sulpharsenate: 

3(NH4)2S  +  As2S3-»2(NH4)3.AsS3, 
3(NH4)2S  +  As2S3  +  2S->2(NH4)3.AsS4. 

Proustite  (p.  417)  is  a  natural  sulpharsenite  of  silver. 

These  salts  are  decomposed  by  acids,  and  give  the  feebly  ionized 
sulpharsenious  or  sulpharsenic  acid: 

(NH4)3.AsS3  +  3HC1  ->  3NH4C1  +  H3AsS3  ->  3H2S  f  +  As2S3  j 
(NH4)3.AsS4  +  3HC1  -»  3NH4C1  +  H3AsS4  ->  3H2S  t  +  As2S5  j. 

These  sulpho-acids,  however,  at  once  break  up,  giving  hydrogen 
sulphide  as  a  gas,  and  the  sulphides  of  arsenic  as  yellow  precipitates. 


ANTIMONY  467 

ANTIMONY  Sb. 

The  Chemical  Relations  of  the  Element.  —  Antimony 
resembles  arsenic  in  forming  a  hydride  SbH3  and  halides  of  the  forms 
SbX3  and  SbX5.  The  latter  are  partially  hydrolyzed  by  water  with 
ease,  but  complete  hydrolysis  is  difficult  to  accomplish  with  cold 
water.  The  oxide  Sb203  is  basic  and  also  feebly  acidic,  and  the 
oxide  Sb2O5  is  acidic.  The  compositions  of  the  compounds  are  similar 
to  those  of  the  compounds  of  arsenic,  but  there  are  in  addition  salts, 
such  as  Sb2(SO4)3,  derived  from  the  oxide  Sb203.  The  element 
gives  complex  sulphides. 

Occurrence  and  Preparation.  —  Antimony  occurs  free  in 
nature.  The  black  trisulphide  Sb2S3,  stibnite,  is  found  in  Hungary 
and  Japan,  and  forms  shining,  prismatic  crystals.  Stibnite  is 
roasted  in  the  air  in  order  to  remove  the  sulphur,  and  the  white 
oxide  which  remains  is  mixed  with  carbon  and  reduced  by  strong 

heat :  Sb2S3  +  502  ->  Sb204  +  3S02, 

Sb204  +  4C  -»  2Sb  +  4CO. 

Properties.  —  Antimony  is  a  white,  crystalline  metal.  It  is 
brittle,  and  easily  powdered.  Its  vapor  at  1640°  has  the  formula 
Sb2,  while  at  lower  temperatures  Sb4  is  present.  It  is  used  in  making 
alloys  such  as  type-metal,  stereotype-metal,  and  britannia  mejal 
(q.v.).  The  alloys  of  antimony  expand  during  solidification,  and 
therefore  give  exceptionally  sharp  castings. 

The  element  unites  directly  with  the  halogens.  It  does  not  rust, 
but  when  heated  it  burns  in  the  air,  forming  the  trioxide  Sb203  or  a 
higher  oxide  Sb204.  When  heated  with  nitric  acid,  it  yields  the 
trioxide  and,  with  more  difficulty,  antimonic  acid  (H3Sb04). 

Stibine  SbHz.  —  The  hydride  of  antimony  SbH3  is  formed  by  the 
action  of  zinc  and  hydrochloric  acid  on  any  soluble  compound  of  anti- 
mony. By  the  action  of  dilute,  cold  hydrochloric  acid  on  an  alloy  of 
antimony  and  magnesium  (1  :  1),  a  mixture  of  hydrogen  and  stibine 
containing  as  much  as  11.5  per  cent  (by  volume)  of  the  latter  may  be 
made.  It  is  separated  by  cooling  with  liquid  air  (b.-p.  -  18°,  m.-p. 
-  91.5°).  It  is  more  easily  dissociated  than  is  arsine  (p.  464),  and 
forms  a  deposit  of  antimony  when  a  porcelain  vessel  is  held  in  the 
flame. 


468  COLLEGE   CHEMISTRY 

Antimony  Halides.  —  The  halides  include  the  trichloride;  the 
pentachloride  SbCl5,  a  liquid  (b.-p.  140°)  ;  the  tribromide  SbBr3,  tri- 
iodide  SbI3,  trifluoride  SbF3,  and  pentafluoride  SbF5. 

Antimony  trichloride  SbCl3  is  made  by  direct  union  of  chlorine  and 
antimony.     It  forms  large,  soft  crystals  (m.-p.  73°),  and  used  to  be 
named  "  butter  of  antimony."     When  treated  with  little  water,  it 
forms  a  white,  opaque,  insoluble  basic  salt,  antimony  oxy  chloride: 
SbCl   +  H0^±SbOCi     +  2HC1. 


With  a  large  amount  of  water,  a  greater  proportion  of  the  chlorine  is 
removed,  and  Sb4O5Cl2  (=  2SbOCl,Sb2O3)  remains.  With  boiling 
water  the  oxide  is  finally  formed.  The  action  is  not  complete  as  long 
as  hydrochloric  acid  is  present.  It  may  therefore  be  reversed,  so 
that,  on  addition  of  hydrochloric  acid  to  the  mixture,  a  clear  solution 
of  the  trichloride  is  re-formed.  If  the  concentration  of  the  acid  is 
once  more  reduced  by  dilution  with  water,  the  oxychloride  is  again 
precipitated. 

Oxides  of  Antimony.  —  The  trioxide  Sb2O3  is  obtained  by 
oxidizing  antimony  with  nitric  acid,  or  by  combustion  of  antimony 
with  a  limited  supply  of  oxygen.  It  is  a  white  substance,  insoluble 
in  water.  It  is  in  the  main  a  basic  oxide,  interacting  with  many 
acids  to  form  salts  of  antimony.  But  it  interacts  also  with  alkalies, 
giving  soluble  antimonites.  The  pentoxide  Sb2O5  is  a  yellow,  amor- 
phous substance,  obtained  by  heating  antimonic  acid.  It  combines 
only  with  bases  to  form  salts,  and  is  therefore  an  acid-forming 
oxide  exclusively.  The  tetroxide  Sb2O4  is  formed  by  heating  anti- 
mony or  the  trioxide  in  excess  of  oxygen.  It  is  neither  acid-  nor 
base-forming. 

Salts  of  Antimony*  —  The  nitrate  Sb(N03)3  and  the  sulphate 
Sb2(S04)3  are  made  by  the  interaction  of  the  trioxide  with  nitric  and 
sulphuric  acids.  They  are  hydrolyzed  by  water,  giving  basic  salts, 
such  as  (SbO)2SO4  (=  Sb2O2SO4),  which,  like  SbOCl,  are  derived  from 
the  hydroxide  SbO(OH).  When  the  trioxide  is  heated  with  a  solu- 
tion of  potassium  bitartrate  KHC4H4O6,  a  basic  salt  K(SbO)C4H4Oe, 
known  as  tartar-emetic,  is  formed.  This  is  a  white,  crystalline  sub- 
stance which  is  soluble  in  water  and  is  used  in  medicine.  The  uni- 
valent  group  SbO1  is  known  as  antimonyl,  and  the  above  mentioned 
basic  compounds  are  often  called  antimonyl  sulphate,  etc. 


ANTIMONY  469 

Antimonic  Acid.  —  By  vigorous  oxidation  of  antimony  with 
nitric  acid,  or  by  decomposing  the  pentachloride  with  water,  a 
white,  insoluble  substance  of  the  approximate  composition  H3SbO4 
is  obtained.  This  substance  interacts  with  caustic  potash  and 
passes  into  solution.  But  the  salts  which  have  been  made  are  pyro- 
and  metantimoniates.  Thus,  when  antimony  is  fused  with  niter, 
potassium  metantimoniate  KSb03  is  formed.  When  dissolved,  this 
salt  takes  up  water,  giving  a  solution  of  the  acid  potassium  pyroanti- 

moniate :  2KSbO3  +  H2O  ->  K2H2Sb207. 

If  this  is  added  to  a  strong  solution  of  a  salt  of  sodium,  an  acid 
sodium  pyroantimoniate  is  thrown  down,  Na2H2Sb2O7.  This  is  almost 
the  only  insoluble  salt  of  sodium. 

Sulphides  of  Antimony.  —  The  trisulphide  Sb2S3  is  found  in 
nature  as  the  black,  crystalline  stibnite.  As  precipitated  from  solu- 
tions of  salts  of  antimony,  the  trisulphide  is  an  orange-red  powder, 
which,  however,  after  melting,  assumes  the  appearance  of  stibnite: 

2SbCl3  +  3H2S  <=±  Sb2S3  J  +  6HC1. 

Antimony  trisulphide,  like  cadmium  sulphide  (p.  435),  cannot  be 
precipitated  in  presence  of  concentrated  hydrochloric  acid. 

The  pentasulphide  Sb2S5  is  obtained  by  the  decomposition,  of 
sulphantimoniates  (see  below).  In  appearance  it  resembles  the 
trisulphide  and,  when  heated,  decomposes  into  this  substance  and 
free  sulphur. 

The  sulphides  of  antimony  behave  towards' solutions  of  the  alkali 
sulphides  as  do  the  sulphides  of  arsenic  (p.  466).  The  trisulphide 
dissolves  in  colorless  ammonium  sulphide  with  difficulty,  forming  an 
unstable,  soluble  ammonium  sulphantimonite 

Sb2S3  +  3(NH4)2S-+2(NH4)3SbS3. 

With  the  pentasulphide  or  with  yellow  ammonium  sulphide  the 
soluble  ammonium  sulphantimoniate  is  readily  formed: 

Sb2S5  +  3(NH4)2S->2(NH4)3.SbS4, 
Sb2S3  +  3(NH4)2S  +  2S->2(NH4)3.SbS4. 

The  most  familiar  substance  of  this  class  is  Schlippe's  salt  Na3SbS4, 
9H20.  Pyrargyrite  (p.  417)  is  a  natural  sulphantimonite. 


470  COLLEGE   CHEMISTRY 

When  acids  are  added  to  solutions  of  sulphantimoniates,  the 
sulphantimonic  acid  which  is  liberated  decomposes,  and  antimony 
pentasulphide  is  thrown  down  (see  under  Arsenic,  p.  466). 

BISMUTH. 

The  Chemical  Relations  of  the  Element.  —  Bismuth  forms 
no  compound  with  hydrogen.  Its  compounds  with  the  halogens  are 
of  the  form  BiX3  and  are  hydrolyzed  by  water  giving  basic  salts. 
The  oxide  Bi2O3  is  basic,  and  the  oxide  Bi2O5  is  not  acidic.  Bismuth 
gives  a  carbonate,  nitrate,  phosphate,  and  other  salts,  in  which  it 
acts  as  a  trivalent  element.  It  forms  no  soluble  complex  sulphides. 

Occurrence  and  Properties.  —  This  element  is  found  free  in 
nature,  and  also  as  trioxide  Bi2O3  and  trisulphide  Bi2S3.  It  is  a 
shining,  brittle  metal  with  a  reddish  tinge  (m.\p.  270°).  Mixtures 
of  bismuth  with  other  metals  of  low  melting-point  fuse  at  lower 
temperatures  than  do  the  separate  metals.  This  is  a  corollary  of 
the  fact  that  a  solution  freezes  at  a  lower  temperature  than  does  the 
pure  solvent  (p.  204).  Thus,  Wood's  metal,  containing  bismuth 
(m.-p.  270°)  4  parts,  lead  (m.-p.  326°)  2  parts,  tin  (m.-p.  233°)  1 
part,  and  cadmium  (m.-p.  320°)  1  part,  melts  at  60.5°,  considerably 
below  the  boiling-point  of  water.  Similar  alloys  are  used  for  safety 
plugs  in  steam-boilers  and  sprinklers. 

Bismuth  does  not  tarnish,  but  when  heated  strongly  it  burns  to 
form  the  trioxide.  With  the  halogens  it  forms  a  fluoride  BiF3,  a 
bromide  BiBr3,  and  an  iodide  BiI3.  When  the  metal  is  treated  with 
oxygen  acids,  or  the  trioxide  with  any  acids,  salts  are  produced. 

Compounds  of  Bismuth.  —  In  addition  to  the  basic  trioxide 
Bi2O3,  which  is  a  yellow  powder  obtained  by  direct  oxidation  of  the 
metal  or  by  ignition  of  the  nitrate,  three  other  oxides  are  known  - 
BiO,  Bi2O4,  and  Bi205.     None  of  these,  however,  is  either  acid- 
forming  or  base-forming. 

The  salts  of  bismuth,  when  dissolved  in  water,  give  insoluble  basic 
salts,  and  the  actions  are  reversible,  the  basic  salts  being  redissolved 
by  addition  of  an  excess  of  the  acid.  In  the  case  of  the  chloride 
BiCl3,H2O  and  the  nitrate  Bi(NO3)3,5H2O,  the  actions  taking  place 


are: 


BiCl3  +  2H2O  +±  Bi(OH)2Cl  +  2HC1, 
Bi(N03)3  +  2H2O^Bi(OH)2N03  +  2HN03. 


THE  FAMILY  AS  A  WHOLE  471 

The  former  of  these  products,  when  dried,  loses  a  molecule  of  water, 
giving  the  oxychloride  BiOCl.  The  oxynitrate  Bi(OH)2N03  is  much 
used  in  medicine  under  the  name  of  "  subnitrate  of  bismuth." 

The  brownish-black  trisulphide  Bi2S3  may  be  obtained  by  direct 
union  of  the  elements,  or  by  precipitation  with  hydrogen  sulphide. 
This  sulphide  is  not  affected  by  solutions  of  ammonium  sulphide  or 
of  potassium  sulphide.  It  differs,  therefore,  markedly  from  the 
sulphides  of  arsenic  and  antimony  in  its  behavior. 

THE  FAMILY  AS  A  WHOLE. 

The  elements  themselves  change  progressively  in  physical  proper- 
ties as  the,  atomic  weight  increases.  Nitrogen  is  a  gas  which  with 
sufficient  cooling  yields  a  white  solid,  phosphorus  an  almost  white, 
or  a  red  solid,  and  arsenic,  antimony,  and  bismuth  are  metallic  in 
appearance.  The  first  combines  directly  with  hydrogen,  the  next 
three  give  hydrides  indirectly,  and  the  last  does  not  unite  with 
hydrogen  at  all.  The  hydride  of  nitrogen  combines  with  water  to 
form  a  base,  while  the  other  hydrides  show  no  such  tendency. 
Ammonia  unites  with  acids,  including  those  of  the  halogens,  to 
form  salts;  phosphine  with  the  hydrogen  halides  only;  the  others  do 
not  combine  with  acids  at  all.  As  regards  their  metallic  properties, 
in  the  chemical  sense,  nitrogen  and  phosphorus  do  not  by  themselves 
form  positive  ions,  and  furnish  us  therefore  with  no  salts  whatever. 
Arsenic  gives  a  trivalent  positive  ion,  which  is  found  in  solutions  of 
the  halides  only.  It  forms  no  normal  sulphates,  nitrates,  or  other 
salts.  Antimony  and  bismuth  both  give  trivalent  positive  ions. 
The  sulphates,  nitrates,  etc.,  of  antimony,  however,  are  readily 
decomposed  by  water  with  precipitation  of  the  hydroxide.  The 
salts  of  bismuth,  on  the  other  hand,  do  not  readily  give  the  pure 
hydroxide  with  water,  although  they  are  easily  hydrolyzed  to  basic 
salts. 

The  halogen  compounds  of  nitrogen  and  phosphorus  are  com- 
pletely hydrolyzed  by  water,  and  do  not  persist  when  any  water  is 
present,  even  when  excess  of  the  halogen  acid  is  used.  The  halogen 
compounds  of  arsenic  are  completely  hydrolyzed  by  cold  water,  but 
exist  in  solution  in  presence  of  excess  of  the  acids.  The  halogen 
compounds  of  antimony  and  bismuth  are  incompletely  hydrolyzed 
by  cold  water. 


472  COLLEGE   CHEMISTRY 

Each  element  gives  a  trioxide  and  a  pentoxide.  With  nitrogen 
these  are  acid-forming,  being  the  anhydrides  of  nitrous  and  nitric 
acids.  With  phosphorus  the  trioxide  and  the  pentoxide  are  an- 
hydrides of  acids.  With  arsenic  the  trioxide  is  basic  towards  the 
halogen  acids,  and  is  the  first  example  of  a  basic  oxide  which  we 
encounter  in  this  group.  The  pentoxide,  however,  is  acid-forming. 
The  trioxide  of  antimony  is  mainly  base-forming,  although  it  is 
feebly  acid-forming  also.  The  pentoxide  is  acid-forming.  The 
trioxide  of  bismuth  is  base-forming  exclusively,  and  the  pentoxide 
has  no  derivatives. 

These  statements,  which  could  easily  be  expanded,  are  sufficient 
to  show  that  when  the  periodic  law  is  borne  in  mind  it  furnishes 
valuable  aid  in  systematizing  the  chemistry  of  a  group  like  this. 

Analytical  Reactions  of  Arsenic,  Antimony,  and  Bismuth. 

—  The  ions  which  are  most  frequently  encountered  are  As"",  Sb*", 
Bi'",  AsO/",  and  AsO/".  The  first  three,  with  hydrogen  sulphide, 
give  colored  sulphides  which  are  not  affected  by  dilute  acids.  The 
sulphides  of  arsenic  and  antimony  are  separable  from  the  sulphide  of 
bismuth  by  solution  in  yellow  ammonium  sulphide.  Marsh's  test 
enables  us  to  recognize  the  presence  of  traces  of  compounds  of 
arsenic  and  antimony.  Oxygen  compounds  of  arsenic,  when  heated 
with  carbon,  give  a  volatile,  metallic-looking  deposit  of  arsenic. 

VANADIUM,  COLUMBIUM,  TANTALUM. 

Of  these  elements,  vanadium  is  less  uncommon  than  the  others. 
It  is  found  in  rather  complex  compounds.  When  these  are  heated 
with  soda  and  sodium  nitrate,  sodium  metavanadate  NaVO3  is  formed, 
and  can  be  extracted  with  water.  The  element  forms  several 
chlorides,  such  as  VC12,  VC13,  VC14,  VOC13,  and  five  oxides,  V20,  VO, 
V203,  VO2,  and  V205.  The  element  has  very  feeble  base-forming 
properties,  and  gives  only  a  few  unstable  salts. 

Columbium  (or  niobium)  and  tantalum  possess  feeble  base-forming 
properties,  their  chief  compounds  being  the  columbates  and  tan- 
talates. 

Exercises.  —  1.  How  do  you  account  for  the  fact  that  the  molec- 
ular weight  of  arsenic  at  644°  is  not  exactly  300,  and  why  is  308.4  -*-  4 
not  accepted  as  the  atomic  weight? 


ARSENIC,  ANTIMONY,  BISMUTH  473 

2.  Formulate  the  series  of  changes  involved  in  the  solution  of 
arsenic  trioxide  and  the  interaction  of  hydrochloric  acid  with  the 
arsenious  acid  so  formed  (cf.  p.  251). 

3.  What  is  the  full  significance  of  the  fact  that  arsenic  penta- 
sulphide  may  be  precipitated  by  hydrogen  sulphide  from  a  solution  of 
arsenic  acid  in  hydrochloric  acid?    Make  the  equation. 

4.  To  what  classes  of  chemical  changes  do  the  interactions  of 
arsenious  sulphide  and  antimony  trisulphide  with  yellow  ammonium 
sulphide  belong? 

5.  Construct  equations  showing  the  interaction  of,  (a)  oxygen  and 
arsenical  pyrites,  (6)  chlorine-water  and  arsenic,  (c)  the  dehydration 
of  orthoarsenic  acid,  (d)  potassium  hydroxide  and  arsenic  trioxide, 
(e)  concentrated  nitric  acid  and  antimony,  (/)  potassium  bitartrate 
and  antimony  trioxide,  (g)  acids  and  ammonium  orthosulphanti- 
moniate. 

6.  How  should  you  set  about  making  Schlippe's  salt? 


: 


CHAPTER  XLI 
THE    CHROMIUM    FAMILY.     RADIUM 

THE  chromium  (Cr,  at.  wt.  52.1)  family  includes  molybdenum  (Mo, 
at.  wt.  96),  tungsten  (W,  at.  wt.  184),  and  uranium  (U,  at.  wt.  238.5), 
and  occupies  the  seventh  column  of  the  periodic  table  along  with  the 
sulphur  and  selenium  family. 

The  Chemical  Relations  of  the  Family.  —  The  features 
which  are  common  to  the  four  elements  are  also  those  which  affiliate 
them  most  closely  with  their  neighbors  on  the  right  side  of  the  column. 
They  yield  oxides  of  the  forms  CrO3,  Mo03,  W03,  and  U03,  which, 
like  S03,  are  acid  anhydrides,  and  show  the  elements  to  be  sexivalent. 
They  give  also  acids  of  the  form  H2XO4,  such  as  chromic  acid  H2Cr04. 
These  acids  correspond  to  sulphuric  acid,  and  their  salts,  for  ex- 
ample the  chromates,  resemble  the  sulphates. 

Aside  from  the  chromates,  the  first  element  forms  also  two  basic 
hydroxides  Cr(OH)2  and  Cr(OH)3,  from  which  the  numerous  chro- 
mous  (Cr")  and  chromic  (Cr'")  salts  are  derived.  Uranium  is  base- 
forming,  as  well  as  acid-forming.  Molybdenum  and  tungsten  are 
not  base-forming  elements. 

CHROMIUM  Cr. 

The  Chemical  Relations  of  the  Element.  —  Chromium  gives 
four  classes  of  compounds,  and  most  of  them  are  colored  substances 
(Gk.  xp&f"1)  color).  The  chromates  are  derived  from  chromic  acid 
H2CrO4,  which,  however,  is  itself  unstable,  and  leaves  the  anhydride 
Cr03  when  its  solution  is  evaporated.  The  oxide  and  hydroxide  in 
which  the  element  is  trivalent,  namely  Cr203  and  Cr(OH)3,  are 
weakly  basic  and  still  more  weakly  acidic.  Hence  we  have  chromic 
salts  such  as  CrCl3  and  Cr2(SO4)3  which  are  somewhat  hydrolyzed, 
but  no  carbonate,  and  no  sulphide  which  is  stable  in  water.  The 
compounds  in  which  the  same  hydroxide  acts  as  an  acid  are  the 

474 


CHROMIUM  475 

chromites,  and  are  derived  from  the  less  completely  hydrated  form 
of  the  oxide  CrO(OH).  Potassium  chromite  K.CrO2  is  more  easily 
hydrolyzed,  however,  than  is  potassium  zincate  or  potassium  alumi- 
nate.  Finally,  the  chromous  salts  such  as  CrCl2  and  CrS04  corre- 
spond to  chromous  hydroxideCr(OH)2in  which  the  element  is  biva- 
lent. This  hydroxide  is  more  distinctly  basic  than  is  chromic 
hydroxide,  and  forms  a  carbonate  and  sulphide  which  can  be  pre- 
cipitated in  aqueous  solution. 

Occurrence  and  Isolation.  —  Chromium  is  found  chiefly  in 
ferrous  chromite  Fe(CrO2)2,  which  constitutes  the  mineral  chromite, 
and  in  crocoisite  PbCr04,  which  is  chromate  of  lead.  The  metal 
may  be  made  by  reduction  of  the  oxide  with  aluminium  filings  by 
Goldschmidt's  method  (p.  444). 

Physical  and  Chemical  Properties.  —  Chromium  is  steel-gray 
in  color,  very  hard,  and  extremely  infusible.  It  does  not  tarnish, 
but  when  heated  it  burns  in  oxygen,  giving  the  green  chromic  oxide 
Cr2O3.  It  seems  to  exist  in  two  states,  an  active  and  a  passive  one, 
the  relations  of  which  are  still  somewhat  obscure.  A  fragment  which 
has  been  made  by  the  Goldschmidt  method,  or  has  been  dipped  in 
nitric  acid,  is  passive,  and  does  not  displace  hydrogen  from  hydro- 
chloric acid.  When,  however,  the  specimen  is  warmed  with  this 
acid,  it  begins  to  interact,  and  thereafter  behaves  as  if  it  lay  between 
zinc  and  cadmium  in  the  electromotive  series.  If  left  in  the  air,  it 
slowly  becomes  inactive  again. 

Tin  and  iron  with  hydrochloric  acid  form  stannous  and  ferrous 
chloride  respectively,  because  the  higher  chlorides,  if  present,  would 
be  reduced  by  the  active  hydrogen  (p.  302).  Here,  for  the  same 
reason,  chromous  chloride  and  not  chromic  chloride  is  formed: 

Cr  +  2HC1  ->  CrCL,  +  H2    or    Cr  +  2H'  ->  Cr"  +  H2. 

DERIVATIVES  OF  CHROMIC  ACID. 

Potassium  Chromate  K2CrO+.  —  This  and  the  sodium  salt,  or 
rather  the  corresponding  dichromates  (see  below),  are  made  directly 
from  chromite,  and  form  the  starting-point  in  the  preparation  of  the 
other  compounds  of  chromium.  The  finely  powdered  mineral  is 
mixed  with  potash  and  limestone,  and  roasted.  The  lime  is  employed 


476  COLLEGE   CHEMISTRY 

chiefly  to  keep  the  mass  porous  and  accessible  to  the  oxygen  of  the 
air,  the  potassium  compounds  being  easily  fusible: 


4Fe(Cr02)2  +  8K2CO3  +  702  -»  2F62O3  +  8K2Cr04  +  8C0 


2. 


The  iron  is  oxidized  to  ferric  oxide,  and  the  chromium  passes  from 
the  state  of  chromic  oxide  in  the  chromite  (FeO,Cr2O3)  to  that  of 
chromic  anhydride  in  the  potassium  chromate  (K2O,CrO3).  Thus, 
more  insight  is  given  into  the  nature  .of  the  action  by  the  equation  : 


4(FeO,Cr203)  +8(K20;C02)  +7p2->2Fe2O3  +  8(K20,CrO3)  +8CO2. 

The  cinder  is  treated  with  hot  potassium  sulphate  solution.  This 
interacts  with  the  calcium  chromate,  which  is  formed  at  the  same 
time,  giving  insoluble  calcium  sulphate: 

CaCr04  +  K2S04  <=?  CaSO4  4  -j-  K2Cr04. 

The  whole  of  the  potassium  chromate  goes  into  solution. 

Potassium  chromate  is  pale-yellow  in  color,  gives  anhydrous, 
rhombic  crystals  like  those  of  potassium  sulphate,  and  is  very  soluble 
in  water  (61  :  100  at  10°). 

Sodium  chromate  Na2CrO4,10H2O  is  made  by  using  sodium  car- 
bonate in  the  process  just  described. 


The  Dichromates.  —  When  a  solution  of  potassium  sulphate  is 
mixed  with  an  equivalent  amount  of  sulphuric  acid,  potassium  bisul- 
phate  is  obtainable  by  evaporation:  K2SO4  +  H2SO4  — >  2KHS04. 
The  dry  acid  salt,  when  heated,  loses  water  (p.  263),  giving  the 
pyrosulphate  (or  disulphate) :  2KHSO4  +±  K2S2O7  +  H20,  but  the 
latter,  when  redissolved,  returns  to  the  condition  of  acid  sulphate. 
Now,  when  an  acid  is  added  to  a  chromate  we  should  expect  the 
chromic  acid  H2Cr04,  thus  liberated,  to  interact,  giving  an  acid 
chromate  (say,  KHCrO4).  No  acid  chromates  are  known,  however, 
and  instead  of  them,  pyrochromates  or  dichromates  are  produced, 
with  elimination  of  water.  In  other  words,  the  second  of  the  above 
actions  is  not  appreciably  reversible  when  chromates  are  in  question : 

K2Cr04   +  H2S04     ->  (H2Cr04)  +  K2SO4 

K,CrO4  (+  H2Cr04)  -»  K2Cr207  +  H20 

2K2Cr04  +  H2S04    ->  K2Cr2O7  +  H2O  +  K2SO4  (1) 


CHROMIUM  477 

In  terms  of  the  ionic  hypothesis,  S2O7"  is  unstable  in  water,  and 
interacts  with  it,  giving  hydrogen-ion  and  sulphate-ion,  while  Cr2O/' 
is  stable  in  water  and  is  formed  from  the  interaction  of  hydrogen-ion 
and  chromate-ion  : 


S2O7"   +  H2O  fc»  2H*  +  2S04", 

Cr207"  +  H2O  <=;  2H'  +  2CrO4".  (2) 

The  dichromates  of  potassium  and  sodium  are  made  by  adding 
sulphuric  acid  to  the  crude  solution  of  the  chromate  obtained  from 
chromite  (p.  476).  They  crystallize  when  the  liquid  cools,  and  the 
mother-liquor,  containing  the  potassium  sulphate  and  undeposited 
dichromate,  is  used  for  extracting  a  fresh  portion  of  cinder.  As  the 
dichromates  are  much  less  soluble  than  the  chromates,  they  crystallize 
from  less  concentrated  solutions,  and  can  therefore  be  obtained  in 
purer  condition.  For^this  reason  the  extract  is  always  treated  for 
dichromate. 

Potassium  diehromate  K2Cr207  (or  K2Cr04,Cr03)  crystallizes  in 
-asymmetric  tables  of  orange-red  color.  Its  solubility  in  water  is 
8  :  100  at  10°  and  12.5  :  100  at  20°/  Sodium  dichromate  Na2Cr207, 
2H20  forms  red  crystals  also,  and  its  solubility  is  109  :  100  at  15°. 
This  salt  is  now  cheaper  than  potassium  dichromate,  and  has  largely 
displaced  the  latter  for  commercial  purposes. 

Chemical  Properties  of  the  Dichromates.  —  1.    When  con- 

centrated sulphuric  acid  is  added  to  a  strong  solution  of  a  dichromate 
(or  chromate),  chromic  anhydride  CrO3  separates  in  red  needles: 

Na2Cr2O7  +  H2SO4  ->  Na^O,  +  H20  +  2CrO3. 

2.  Although  a  dichromate  lacks  the  hydrogen,  it  is  essentially  of 
the  nature  of  an  acid  salt,  just  as  SbOCl  lacks  hydroxyl,  but  is  essen- 
tially a  Basic  saItT~  Hence7  when  potassium  hydroxide  is  added  to  a 
solution  of  potassium  dichromate,  potassium  chromate  is  formed  : 

K2Cr207  +  2KOH  -»  2K2CrO4  +  H20. 

The  solution  changes  from  red  to  yellow,  and  the  chromate  is  obtained 
by  evaporation.     In  this  way  the  pure  alkali  chromates  are  made. 

3.  By  addition  of  potassium  dichromate  to  a  solution  of  a  salt  of  a 
metal  whose  chromate  is  insoluble,  the  chromate  and  not  the  dichro- 
mate is  precipitated.     This  is  in  consequence  of  the  fact  that  there  is 


478  COLLEGE   CHEMISTRY 

always  a  little  hydrogen-ion  and  O04"  (equation  (2),  above)  in  the 
solution  of  the  dichromate: 

2Ba(NO3)2  +  K2Cr2O7  +  H2O  <=»  2BaCrO4  J  +  2KNO3  +  2HN03. 

Being  essentially  an  acid  salt,  the  dichromate  produces  a  salt  and  an 
acid,  as  any  acid  salt  would  do.     For  example: 

Ba(NO3)2  +  KHSO4  «=>  BaSO4  1  +  KNO3  +  HNO3. 

4.  The  dichromates  of  potassium  and  sodium  melt  when  heated 
and,  at  a  white  heat,  decompose,  giving  the  chromate,  chromic  oxide, 
and  free  oxygen.     To  make  the  equation,  we  note  that  the  dichro- 
mate, for  example  K2O2O7,  consists  of  K2CrO4  -f  Cr03,  and  the 
latter,  if  alone,  will  decompose  thus:  2CrO3  —  >  Cr2O3  +  30.     Since 
the  product  must  contain  a  multiple  of  O2,  the  equation  is: 

4K2O2O7  ~»  4K2CrO4  +  2Cr2O3  +  302. 

5.  With  free  acids  the  dichromates  give  powerful  oxidizing  mix- 
tures, in  consequence  of  their  tendency  to  form  chromic  salts.     Since 
the  latter  correspond  to  the  oxide  Cr2O3  and  the  former  to  CrO3,  the 
passage  from  the  former  to  the  latter  must  furnish  3O  for  every 
2Cr03  transformed.     In  dilute  solutions,  unless  a  body  capable  of 
being  oxidized  is  present,  no  actual  decomposition,  beyond  the  libera- 
tion of  chromic  acid  *  occurs.     When  concentrated  hydrochloric 
acid  is  used  this  acid  itself  suffers  oxidation: 


K2Cr2O7  4 
(30)        4 

-    8HC1  ->  2KC1  H 
6HC1  -»  3H20  H 

-  2CrCl3  - 
h3C!2 

1-  4H2O  (  +  3 

K2Cr2O7  +  14HC1  -*  2KC1  +  2CrCl3  +  7H2O  +  3C12 

When  sulphuric  acid  is  employed,  an  oxidizable  substance  such  as 
hydrogen  sulphide  (cf.  p.  253)  ,  sulphurous  acid,  or  alcohol  must  be 
present,  if  the  dichromate  is  to  be  reduced  : 

K2Cr207  +     4H2SO4  -*  K2SO4  +  Cr2(SO4)3  +  4H2O(+  30)    (1) 
(3O)  +      3H2SO3  -4  3H2S04  (2) 

or         (30)  +  3C2H5OH  -*  3C?H4O  |  +  3H2O  (20 

[alcohol]  [aldehyde] 

In  each  case  the  usual  summation  of  (1)  and  (2),  with  omission  of  the 
3O  gives  the  equation  for  the  whole  action.  When  (1)  is  dissected, 

*  Not  shown  as  a  distinct  stage  in  the  subsequent  equations. 


CHROMIUM  479 

K20,2CrO3  giving  Cr203,3S03  +  3O  is  found  to  be  its  essential 
content.  In  practice,  this  sort  of  action  is  used  for  the  purpose  of 
making  chromic  salts,  and  for  its  oxidizing  effects,  as  in  the  prepara- 
tion of  aldehyde  and  in  the  dichromate  battery. 

6.  When  paper  is  coated  with  gelatine  containing  a  soluble 
chromate  or  dichromate,  and,  after  being  dried,  is  exposed  to  light, 
chromic  oxide  is  formed  by  reduction,  and  combines  with  the  gela- 
tine. This  product  will  not  swell  up  or  dissolve  in  tepid  water,  as 
does  pure  gelatine.  This  action  is  used  in  many  ways  for  purposes 
of  artistic  reproduction.  Thus,  if  the  gelatine  mixture  is  made  up 
with  lampblack,  and,  after  the  coating  has  dried,  is  covered  with  a 
negative  and  exposed  to  light,  the  parts  which  were  protected  from 
illumination  may  afterwards  be  washed  away,  while  the  carbon  print 
remains.  The  gelatine  layer  can  be  transferred  to  wood  or  copper 
before  washing.  When  materials  of  different  colors  are  substituted 
for  the  lampblack,  prints  of  any  desired  tint  may  be  made  by  the 
same  process. 

Insoluble  Chromates.  —  A  number  of  chromates,  formed  by  pre- 
cipitation with  a  solution  of  a  soluble  chromate  or  dichromate,  are 
familiar.  Thus,  lead  chromate  PbCrO4  is  used  as  a  yellow  pigment. 
By  treatment  with  lime-water  it  gives  a  basic  salt  of  brilliant  orange 
color,  —  chrome-red  Pb2OCrO4.  Salts  of  calcium  give  a  yellow, 
hydrated  calcium  chromate  CaCrO4,  2HO2  analogous  to  gypsum,  and, 
like  it,  perceptibly  soluble  in  water  (0.4  :  100  at  14°).  Barium  chro- 
mate BaCrO4  is  also  yellow.  Being  a  salt  of  a  feeble,  acid,  it  interacts 
with  active  acids,  and  passes  into  solution.  Like  calcium  oxalate 
(cf.  p.  397),  it  is  not  soluble  enough  to  be  attacked  by  acetic  acid. 
Strontium  chromate  SrCr04,  however,  is  soluble  in  acetic  acid.  Silver 
chromate  Ag2CrO4  is  red,  and  interacts  easily  with  acids.  It  will  be 
observed  that  there  is  a  close  correspondence  between  the  relative 
'solubilities  (see  Table)  of  the  chromates  and  the  sulphates. 

Chromic  Anhydride  OO3.  —  This  oxide  is  made  as  described 
above  (par.  1,  p.  477),  and  is  often  called  chromic  acid.  It  is  soluble 
in  water,  and  combines  with  the  latter  to  some  extent,  giving  di- 
chromic acid  H2.O207.  In  a  solution  acidified  with  an  active  acid  it 
is  much  used  as  an  oxidizing  agent  for  organic  substances.  It  inter- 
acts with  acids  in  the  same  way  as  do  the  dichromates,  giving  chromic 


480  COLLEGE   CHEMISTRY 

salts  and  furnishing  oxygen  to  the  oxidizable  body.  When  heated  by 
itself,  it  loses  oxygen  readily,  and  yields  the  green  chromic  oxide: 
4Cr03-»2Cr203  +  302. 

CHROMIC  AND  CHROMOUS  COMPOUNDS. 

Chromic  Chloride.  —  A  hydrated  chloride  CrCl3,  6H20  is  ob- 
tained by  treating  the  hydroxide  Cr(OH)3  with  hydrochloric  acid 
and  evaporating.  When  heated,  this  hydrate  is  hydrolyzed,  and 
chromic  oxide  remains.  The  anhydrous  chloride  is  formed  by  subli- 
mation, as  a  mass  of  brilliant,  reddish-violet  scales,  when  chlorine  is 
led  over  a  heated  mixture  of  chromic  oxide  and  carbon  (cf.  p.  344) : 
Cr203  +  3C  +  3C12  ->  2CrCl3  +  SCO. 

In  this  form  the  substance  dissolves  with  extreme  slowness,  even  in 
boiling  water,  but  in  presence  of  a  trace  of  chromous  chloride  or 
stannous  chloride  it  is  easily  soluble.  The  solution  is  green,  as  are 
all  solutions  of  chromic  salts  after  they  have  been  boiled,  but  on 
standing  in  the  cold,  bluish  crystals  of  CrCl3,  6H2O  are  deposited. 
These  give  a  bluish  solution  containing  Cr""  +3  Cl',  but  boiling  repro- 
duces the  green  color.  The  green  material  is  a  basic  salt  of  an  ex- 
ceptional nature.  It  has  lost  one  unit  of  chlorine  by  hydrolysis,  and 
one  of  the  two  others  is  not  precipitated  by  nitrate  of  silver.  The 
substance  is  supposed,  therefore,  to  contain  a  complex  cation  and  to 
have  the  formula  CrClOH.Cl. 

Chromic  Hydroxide.  —  When  ammonium  hydroxide  is  added  to 
a  solution  of  a  chromic  salt,  a  hydrated  hydroxide  of  pale-blue  color, 
Cr(OH)3,2H2O,  is  thrown  down.  This  loses  water  by  stages,  giving 
intermediate  hydroxides  such  as  Cr(OH)3  and  CrOOH,  and  finally 
Cr203.  It  interacts  with  acids,  giving  chromic  salts.  It  also  dis- 
solves in  potassium  and  sodium  hydroxides  to  form  green  solutions 
of  chromites  of  the  form  KCrO2.  When  the  solutions  of  the  alkali 
chromites  are  boiled,  the  free  chromic  hydroxide,  present  in  conse- 
quence of  hydrolysis,  is  converted  into  a  greenish,  less  completely 
hydrated,  and  less  soluble  variety.  This  begins  to  come  out  as  a  pre- 
cipitate, and  soon  the  whole  action  is  reversed.  Insoluble  chromites, 
such  as  that  of  ironFe(CrO2)2,  are  found  in  nature.  Many  of  them,  like 
Zn(CrO2)2  and  Mg(Cr02)2,  may  be  formed  by  fusing  the  oxide  of  the 
metal  with  chromic  oxide;  the  action  being  similar  to  that  used  in 
making  zincates  (p.  433)  and  aluminates  (p.  450). 


CHROMIUM  481 

Chromic  Oxide  Cr2O3.  —  This  oxide  is  obtained  as  a  green, 
infusible  powder  by  heating  the  hydroxide;  or,  more  readily,  by 
heating  dry  ammonium  dichromate;  or  by  igniting  potassium 
dichromate  with  sulphur  and  washing  the  potassium  sulphate  out 
of  the  residue: 

(NH4)20207  -»  N2  +  4H20  +  Cr203, 
K2O2O7  +  S  ->  K2SO4  +  Cr2O3. 

Chromic  oxide  is  not  affected  by  acids,  but  may  be  converted  into 
the  sulphate  by  fusion  with  potassium  bisulphate.  It  is  used  for 
making  green  paint,  and  for  giving  a  green  tint  to  glass.  When  the 
oxide,  or  any  of  the  chromic  salts,  is  fused  with  a  basic  substance  such 
as  an  alkali  carbonate,  it  passes  into  the  form  of  a  chromate,  absorb- 
ing the  necessary  oxygen  from  the  air.  If  an  alkali  nitrate  or 
chlorate  is  added,  the  oxidation  goes  on  more  quickly.  The  alkaline 
solution  of  the  chromites  may  be  oxidized,  for  example  by  chlorine 
or  bromine,  and  chromates  are  formed. 

Chromic  Sulphate  Cr2(SO4)3, 15H2O.  —  This  salt  crystallizes  in 
reddish-violet  crystals,  and  may  be  made  by  treating  the  hydroxide 
with  sulphuric  acid.  It  gives  reddish-violet,  octahedral  crystals  of 
chrome-alum  (cf.  p.  447),  K2S04,  Cr2(S04)3,  24H20,  when  mixed  with 
potassium  sulphate.  This  double  salt  is  most  easily  obtained  by 
reducing  potassium  dichromate  in  dilute  sulphuric  acid  by  means  of 
sulphurous  acid  (p.  478),  and  allowing  the  solution  to  crystallize. 
The  solution  of  the  crystals,  either  of  the  pure  sulphate  or  of  the  alum, 
is  bluish-violet  (Cr***),  but  when  boiled  becomes  green.  The  green 
compound  is  formed  by  hydrolysis  and  is  gummy  and  uncrystal- 
lizable.  It  even  yields  products  which  do  not  show  the  presence 
either  of  the  Cr"*  or  the  S04"  ion.  It  seems  to  be  formed  thus : 

2Cr2(S04)3  +  H20<=>Cr40(S04)4.S04  +  H2SO4. 

The  green  materials  revert  slowly  to  the  violet  ones  by  reversal  of 
the  above  action  when  the  solution  remains  in  the  cold,  and  so 
crystals  of  the  sulphate  or  of  the  alum  are  obtainable  from  the  green 
solutions. 

Chromous  Compounds.  —  By  the  interaction  of  chromium  with 
hydrochloric  acid,  or  by  reducing  chromic  chloride  in  a  stream  of 
hydrogen,  chromous  chloride  CrCl2  is  formed.  The  anhydrous  salt  is 


482  COLLEGE   CHEMISTRY 

colorless,  and  its  solution  is  blue  (Cr").  Like  stannous  chloride,  it 
is  very  easily  oxidized  by  the  air,  a  solution  of  it  containing  excess  of 
hydrochloric  acid  being  used  in  the  laboratory  to  absorb  oxygen: 

4CrCl2  +  4HC1  +  O2  -»  4CrCl3  +  2H,O. 

Chromous  hydroxide  Cr(OH)2  is  obtained  as  a  yellow  precipitate 
when  alkalies  are  added  to  the  chloride.  With  sulphuric  acid  it 
gives  chromous  sulphate  CrSO4,  7H20,  which  is  one  of  the  vitriols 
(p.  433). 

Analytical    Reactions    of    Chromium     Compounds.  —  The 

chromic  salts  give  the  bluish-violet  chromic-ion  Cr"*,  or  the  green 
complex  cations,  and  may  be  recognized  in  solution  by  their  color. 
The  chromates  and  dichromates  give  the  ions  OO4"  and  Cr2O7", 
which  are  yellow  and  red  respectively.  From  chromic  salts,  alkalies 
and  ammonium  sulphide  precipitate  the  bluish-green  hydroxide,  and 
carbonates  give  a  basic  carbonate  which  is  almost  completely 
hydrolyzed  to  hydroxide.  By  fusion  with  sodium  carbonate  and 
sodium  nitrate,  they  yield  a  yellow  bead  containing  the  chromate. 
The  chromates  and  dichromates  are  recognized  by  the  insoluble 
chromates  which  they  precipitate,  and  by  their  oxidizing  power 
when  mixed  with  acids.  All  compounds  of  chromium  give  a 
green  borax  bead  containing  chromic  borate,  and  this  bead 
differs  from  that  given  by  compounds  of  copper  (cf.  p.  417),  both  in 
color  and  in  being  unreducible. 

MOLYBDENUM,  TUNGSTEN,  URANIUM. 

Molybdenum.  —  This  element  is  found  chiefly  in  wulfenite 
PbMo04  and  molybdenite  MoS2.  The  latter  resembles  black  lead 
(graphite),  and  its  appearance  suggested  the  name  of  the  element 
(Gk.  /noAv'/3oWa,  lead).  The  molybdenite  is  converted  by  roasting 
into  molybdic  anhydride  MoO3.  When  this  is  treated  with  ammonium 
hydroxide,  or  with  sodium  hydroxide,  ammonium  molybdate  (NH4)2 
MoO4,  or  sodium  molybdate  Na-jMoO,,,  10H20  is  obtained.  The  metal 
itself  is  liberated  by  reducing  the  oxide  or  chloride  with  hydrogen. 
When  pure  it  resembles  wrought  iron  and,  like  iron  (q.v.),  takes  up 
carbon  and  shows  the  phenomena  of  tempering.  The  oxides  [MoO  ?], 
Mo2O3,  Mo02,  and  MoO3  are  known,  but  the  lower  oxides  are  not 


MOLYBDENUM,    TUNGSTEN,   URANIUM  483 

basic.  The  chlorides  Mo3Cle,  MoCl3,  MoCl4,  and  MoCl5  have  been 
made.  The  chief  use  of  molybdenum  compounds  in  the  laboratory 
is  in  testing  for  and  estimating  phosphoric  acid.  When  a  little  of  a 
phosphate  is  added  to  a  solution  of  ammonium  molybdate  in  nitric 
acid,  and  the  mixture  is  warmed,  a  copious  yellow  precipitate  of  a 
phosphomolybdate  of  ammonium  (NH4)3PO4,llMoO3,6H20  is  formed. 
The  compound  is  soluble  in  excess  of  phosphoric  acid  and  in  alkalies, 
but  not  in  dilute  mineral  acids. 


Tungsten.  —  The  minerals  scheelite  CaWO4  and  wolfram  FeW04 
are  tungstates  of  calcium  and  iron  respectively.  By  fusion  of 
wolfram  with  sodium  carbonate  and  extraction  with  water,  sodium 
tungstate  Na2WO4,2H2O  is  secured.  It  is  used  as  a  mordant  and  for 
rendering  muslin  fire-proof.  Acids  precipitate  tungstic  acid  H2WO4, 
H2O  from  solutions  of  this  salt.  The  element  gives  the  oxides  WO2 
and  WO3,  the  latter  being  formed  by  ignition  of  tungstic  acid.  The 
chlorides  WC12,  WC14,  WC15,  and  WC10  are  known,  the  last  being 
formed  directly,  and  the  others  by  reduction.  A  hard  variety  of 
steel  contains  5  per  cent  of  tungsten. 


Uranium.  —  This  element  is  found  chiefly  in  pitchblende,  which 
contains  the  oxide  U308  along  with  smaller  amounts  of  many  other 
elements.  By  roasting  the  ore  with  carbonate  and  nitrate  of  sodium, 
and  extracting  with  water,  an  impure  solution  of  sodium  uranate 
NaU04  is  obtained.  Acids  precipitate  the  insoluble,  yellow  diuranate 
Na2U207,  6H20.  This  salt  is  used  in  making  uranium  glass,  which 
shows  a  yellowish-green  fluorescence.  The  property  is  due  to  the 
fact  that  the  wave-length  of  part  of  the  invisible,  ultra-violet  rays  of 
the  sunlight  are  shortened,  and  a  greenish  light  is  therefore  in  excess. 
The  oxides  are  U02  a  basic  oxide,  U2O3,  U3O8  the  most  stable  oxide, 
U03  uranic  anhydride,  and  U04  a  peroxide. 

When  the  oxide  U02  is  treated  with  acids,  it  gives  uranous  salts 
such  as  uranous  sulphate  U(S04)2,  4H20.  Uranic  anhydride  and 
uranic  acid  interact  with  acids,  giving  basic  salts,  such  as  UO2S04, 
3JH20,  and  U02(N03)2,  6H20,  which  are  named  uranyl  sulphate, 
uranyl  nitrate,  and  so  forth.  They  are  yellow  in  color,  with  green 
fluorescence.  Ammonium  sulphide  throws  down  the  brown,  unstable 
uranyl  sulphide  U02S  from  their  solutions. 


484  COLLEGE    CHEMISTRY 

RADIUM. 

The  Discovery  of  the  Element.  *  —  It  was  Becquerel  who 
first  noticed  (1896)  that  all  compounds  of  uranium  give  out  a  radia- 
tion capable  of  slowly  affecting  a  photographic  plate  covered  with 
black,  light-proof  paper.  These  rays  have  a  second,  equally  remark- 
able property.  A  well-insulated,  electrically  charged  body,  such  as 
an  electroscope,  will  retain  its  charge  for  a  long  time.  Yet  a  few 
tenths  of  a  gram  of  any  uranium  compound,  brought  within  3  or  4 
cm.  of  the  charged  body,  render  the  air  a  conductor,  and  the  charge 
is  quickly  lost.  The  air  is  said  to  be  "  ionized."  For  the  quantita- 
tive measurement  of  radio-activity,  or  the  rate  of  production  of 
Becquerel  rays,  we  simply  have  to  compare  the  times  required  for 
the  discharge  of  an  electroscope  by  different  specimens  of  radio- 
active matter. 

The  radio-activity  of  every  pure  uranium  compound  is  propor- 
tional to  its  uranium  content.  The  ores  are,  however,  relatively  four 
times  as  active.  This  fact  led  M.  and  Mme.  Curie  to  the  discovery 
that  the  pitchblende  residues,  from  which  practically  all  of  the 
uranium  had  been  extracted,  were  nevertheless  quite  active.  About 
a  ton  of  the  very  complex  residues  having  been  separated  laboriously 
into  the  constituents,  it  was  found  that  a  large  part  of  the  radio- 
activity remained  with  the  chloride  of  barium.  From  this  barium 
chloride,  a  product  free  from  barium,  and  three  million  times  as 
active  as  uranium,  was  finally  secured.  The  nature  of  the  spectrum 
and  the  chemical  relations  of  the  element,  now  named  radium,  placed 
it  with  the  metals  of  the  alkaline  earths.  The  ratio  by  weight  of 
chlorine  to  radium  in  the  compound  is  35.45  :  113.25,  so  that,  on  the 
assumption  that  the  element  is  bivalent,  its  atomic  weight  is  326.5. 
With  this  value  it  occupies  a  place  formerly  vacant  in  the  periodic 
table. 

Properties  of  Radium  Compounds.  —  Radium  has  not  been 
isolated,  and  the  chloride  RaCl2  and  bromide  RaBr2  are  usually 
employed.  The  extraordinary  photo-activity  and  ionizing  power 
of  the  compounds  have  been  mentioned  above.  The  most  remark- 
able property  of  the  salts,  however,  is  their  constant  evolution  of 
heat  in  relatively  enormous  quantities.  One  gram  of  the  element, 

*  I  am  indebted  to  my  colleague  Professor  H.  N.  McCoy  for  the  material  of 
which  the  following  is  a  slightly  condensed  version. 


RADIUM  485 

in  combination,  evolves  over  100  cal.  per  hour,  and  it  is  estimated 
that  the  total  amount  of  heat  spontaneously  produced  by  1  g.  would 
be  nearly  1010  cal.  To  produce  the  same  amount  of  heat  by  combus- 
tion, no  less  than  300  kg.  of  hydrogen  would  have  to  be  burned.  The 
heating  effect  is  explained  by  the  hypothesis  that  the  rays  emitted 
by  the  radium  consist  of  minute  particles  whose  impacts  produce 
the  heat.  To  explain  all  the  properties  we  are  compelled  to  suppose 
that  there  are  two  kinds  of  these  particles,  electrons,  or  corpuscles, 
namely  those  which  produce  the  ionization  (a-particles)  and  those 
which  produce  the  photographic  effects  (/J-particles).  The  former 
bear  positive  charges  and  have  twice  the  mass  of  a  hydrogen  atom. 
The  latter  bear  negative  charges  and  have  only  about  1/1 000th  of 
the  mass  of  a  hydrogen  atom. 

The  Decay  of  an  Element.  —  The  radium  in  disintegrating 
and  giving  off  these  particles,  produces  a  new  radio-active,  gaseous 
body,  the  radium  emanation.  This  is  not  permanent,  and  loses  half 
its  activity  in  about  four  days.  In  doing  so  it  produces  a  series  of 
new  radio-active  substances.  The  members  of  the  series  with  their 
times  of  decay  to  half  value,  are  as  .follows  (Rutherford) : 

U      — >        U-X  — >  Ionium     — >      Ra      — ->        Em    — >     A     — » 
5Xl09yrs-   22  days  ?  2000  yrs.      3.8  days      3  min. 

B->        C->D->E->F-+        G 
26  min.        19  min.        40  yrs.     6  days     45  days         140  days 

Radium  thus  loses  its  activity.  To  account  for  its  presence  in  the 
ore,  we  must,  therefore,  suppose  that  it  is  being  continuously  pro- 
duced from  some  source.  This  source  must  be  uranium,  for  every 
known  uranium  ore  contains  radium  (McCoy)  and  radium  emanation 
(Boltwood)  in  amounts  proportional  to  the  uranium  content.  Fur- 
thermore, the  radium  emanation  is  slowly  re-formed  in  a  uranium 
compound  previously  freed  from  radium  (Soddy). 

It  has  been  found  (Ramsay  and  Soddy)  that  helium  (p.  290)  is 
one  of  the  decomposition  products  of  the  radium  emanation,  and, 
more  recently,  that  cupric  nitrate  solution,  when  inclosed  with  some 
of  the  radium  emanation,  acquires  a  small  proportion  of  lithium 
nitrate  (Ramsay  and  Cameron).  The  fact  that  all  uranium-radium 
minerals  which  contain  copper,  contain  also  lithium  (McCoy),  con- 
firms this  observation  of  the  degradation  of  copper  to  that  member 


486  COLLEGE   CHEMISTRY 

of  the  same  family  possessing  the  lowest  atomic  weight.  The  phe- 
nomena of  radio-active  substances  lead  undeniably  to  the  startling 
conclusion  that  some,  if  not  all,  of  the  elements  are  capable  of  spon- 
taneous decomposition. 

Exercises.  —  1 .   Construct  equations,  showing  the  interactions  of : 

(a)  chromic  oxide  and  aluminium,  (6)  strontium  nitrate  and  potas- 
sium dichromate  in  solution,  (c)  potassium  hydroxide  and  chromic 
hydroxide,  and  the  reversal  on  boiling,  (d)  chlorine  and  potassium 
chromite  in  excess  of  alkali  (what  is  the  actual  oxidizing  agent?). 

2.  What  volume  of  oxygen  at  0°  and  760  mm.,  (a)  is  obtainable 
from  one  formula-weight  of  potassium  dichromate  (par.  4,  p.  .478), 

(b)  is  required  to  oxidize  one  formula-weight  of  chromous  chloride? 

3.  To  what  classes  of  actions  should  you  assign  the  three  methods 
of  making  chromic  oxide  (p.  481)? 


CHAPTER  XLII 
MANGANESE 

Tfie  Chemical  Relations  of  the  Element.  —  Manganese 
stands,  at  present,  alone  on  the  left  side  of  the  eighth  column  of  the 
periodic  table.  The  right  side  is  occupied  by  the  halogens.  It  is 
never  univalent,  as  the  halogens  are,  but  its  heptoxide  Mn2O7  and 
the  corresponding  acid,  permanganic  acid  HMnO4,  are  in  many  ways 
closely  related  to  the  heptoxide  of  chlorine  and  perchloric  acid  HC1O4. 
Of  the  lower  oxides  of  manganese,  MnO  is  basic,  and  Mn203  feebly 
basic.  Mn02  is  feebly  acidic,  Mn03  more  strongly  so,  and  perman- 
ganic acid  (from  Mn207)  is  a  very  active  acid.  Contrary  to  the  habit 
of  feebly  acidic  and  feebly  basic  oxides,  such  as  those  of  zinc,  alumi- 
nium, and  tin,  the  basic  oxides  of  manganese  are  not  at  all  acidic, 
and  the  acidic  oxides  (with  the  possible  exception  of  Mn2O3)  are  not 
also  basic.  There  are  thus  the  five  following,  rather  well-defined  sets 
of  compounds,  showing  five  different  valences  of  the  element.  Of 
these  the  first,  fourth,  and  fifth  are  the  most  stable  and  the  most 
important. 

1.  Manganous  compounds,  MnO,  Mn(OH)2,  MnS04,  etc.     These 
compounds  resemble  those  of  the  magnesium  family  (and  those  of 
Fe").     The  salts  of  weak  acids,  such  as  the  carbonate  and  sulphide, 
are  easily  made,  and  there  is  little  hydrolysis  of  the  halides.     The 
salts  are  pale-pink  in  color. 

2.  Manganic  compounds,  Mn2O8l   Mn(OH)3,  Mn2(SO4)3,  [MnCl3]. 
The  salts  resemble  the  chromic  and  aluminium  salts  in  behavior,  but 
are  even  less  stable  than  those  of  quadrivalent  lead.    They  are 
completely   hydrolyzed   by  little  water.     The  salts  are  violet  in 
color. 

3.  Manganites,  MnO2,  H2MnO3,  CaMnO3.     The  alkali  manganites 
are  strongly  hydrolyzed,  like  the  plumbates  and  the  stannates. 

4.  Manganates,  Mn03,  HJVInO4,  K2MnO4.     The  salts  resemble  the 
sulphates  and  chromates,  but  are  much  more  easily  hydrolyzed. 
The  free  acid  resembles  chloric  acid  in  that  when  it  decomposes  it 

487 


488  COLLEGE   CHEMISTRY 

yields  a  higher  acid  (HMn04)  and  a  lower  oxide  (MnO2).  The  salts 
are  green  in  color. 

5.  Permanganates,  Mn2O7,  HMnO4(hydrated),  KMn04.  The  salts 
resemble  the  perchlorates,  and  are  not  hydrolyzed  by  water.  They 
are  reddish-purple  in  color. 

It  will  be  seen  that  the  element  manganese  changes  its  character 
totally  with  change  in  valence,  and  in  each  form  of  combination 
resembles  some  set  of  elements  of  valence  identical  with  that  which 
it  has  itself  assumed. 

Occurrence :  the  Metal.  —  The  chief  ore  is  the  dioxide,  pyro- 
lusite  MnO2,  which  always  contains  compounds  of  iron.  Other  man- 
ganese minerals  are:  braunite  Mn2O3;  the  hydrated  form,  manganite 
MnO(OH);  hausmannite  Mn3O4;  and  manganese  spar  MnCO3.  The 
metal  is  most  easily  made  by  reducing  one  of  the  oxides  with  alumi- 
nium by  Goldschmidt's  method. 

The  metal  manganese  has  a  grayish  luster  faintly  tinged  with  red. 
It  rusts  in  moist  air,  and  easily  displaces  hydrogen  from  dilute  acids, 
giving  manganous  salts.  Its  alloys  with  iron,  such  as  ferro-manga- 
nese  (20-80  per  cent  manganese) ,  are  used  in  the  arts. 

Oxides.  —  Manganous  oxide  MnO  is  a  green  powder,  made  by 
reducing  any  of  the  other  oxides  with  hydrogen.  Hausmannite  Mn3O4 
is  red.  An  oxide  having  this  composition  is  formed  when  any  of  the 
other  oxides  is  heated  in  air,  oxidation  or  reduction,  as  the  case  may 
be,  taking  place  (cf.  p.  458).  Manganic  oxide  Mn203  is  brownish- 
black,  and  is  formed  by  heating  any  of  the  oxides  in  oxygen. 

Manganese  dioxide  Mn02  is  black,  and  is  most  easily  prepared  in 
pure  condition  by  gentle  ignition  of  manganous  nitrate.  The 
hydrated  forms  of  the  oxide  are  produced  by  precipitation,  as  by 
adding  a  hypochlorite  or  hypobromite  to  manganous  hydroxide 
suspended  in  water.  Manganese  dioxide  is  not  a  peroxide  in  the 
restricted  sense  (cf.  p.  212).  That  is  to  say,  it  does  not  contain  the 
radical  (02)  and,  therefore,  does  not  give  hydrogen  peroxide.  Its 
reaction  formula  is  Mn(O)2  not  Mn(O,)  and  in  double  decompositions 
it  yields  only  water  H2(O).  It  is  used  for  manufacturing  chlorine, 
although  electrolytic  processes  are  now  driving  it  out  of  this  field. 
In  glass-making  (q.v.),  it  is  employed  to  oxidize  the  green  ferrous 
silicate,  derived  from  impurities  in  the  sand,  to  the  pale-yellow  ferric 


MANGANESE  489 

compound.  The  amethyst  color  of  the  manganic  silicate  which  is 
formed  tends  to  neutralize  this  yellow. 

Manganese  trioxide  MnO8  is  a  red,  unstable  powder.  Manganese 
heptoxide  Mn2O7  is  a  brownish-green  oil  (see  below). 

When  any  of  these  oxides  is  heated  with  an  acid,  a  manganous 
salt  is  obtained.  Salts  of  this  class  are,  in  fact,  the  only  stable  sub- 
stances in  which  manganese  is  combined  with  an  acid  radical.  In 
this  action  the  oxides  containing  more  oxygen  than  does  MnO  give 
off  oxygen,  or  oxidize  the  acid  (cf.  p.  112).  When  the  oxides  are 
heated  with  bases,  in  the  presence  of  air,  manganates  are  always 
formed.  In  this  case,  with  oxides  containing  a  smaller  proportion 
of  oxygen  than  MnO3,  oxygen  is  taken  from  the  air. 

Manganous  Compounds* — The  manganous  salts  are  formed  by 
the  action  of  acids  upon  the  carbonate  or  any  of  the  oxides.  Thus 
the  chloride  MnCl2,4H2O  is  obtained  in  pale-pink  crystals  from  a 
solution  made  by  treating  the  dioxide  with  hydrochloric  acid  and 
driving  off  the  chlorine  liberated  by  oxidation  (p.  112).  The  hydrox- 
ide Mn(OH)2  is  formed  as  a  white  precipitate  when  a  soluble  base  is 
added  to  a  solution  of  a  manganous  salt.  This  body  passes  into 
solution  when  ammonium  salts  are  added,  and  cannot  be  precipitated 
in  their  presence  on  account  of  the  formation  of  molecular  ammonium 
hydroxide  and  the  suppression  of  hydroxide-ion  (cf.  magnesium 
hydroxide,  p.  429).  The  hydroxide  quickly  darkens  when  exposed 
to  the  air  and  passes  over  into  hydrated  manganic  oxide  MnO  (OH). 

Manganous  sulphate  gives  pink  crystals  of  a  hydrate.  Below  6° 
the  solution  deposits  MnSO4,7H2O,  which  is  a  vitriol  (p.  433). 
Between  7°  and  20°  the  product  is  MnSO4,5H20,  asymmetric  and 
resembling  CuSO4,5H2O.  Above  25°  monosymmetric  prisms  of 
MnSO4,4H2O  are  obtained.  These  hydrates  have  different  aqueous 
tensions  and  may  be  formed  from  one  another  by  lowering  or  raising 
the  pressure  of  water  vapor  around  the  substance  (p.  83). 

Manganous  carbonate  MnCO3  is  a  white  powder  formed  by  pre- 
cipitation. The  sulphide  MnS  is  obtained  as  a  green  powder  by 
leading  hydrogen  sulphide  over  any  of  the  oxides.  A  flesh-colored, 
hydrated  manganous  sulphide  MnS,  H2O  or  MnSH(OH)  is  more 
familiar  and  is  precipitated  by  ammonium  sulphide  from  manganous 
salts.  It  interacts  with  mineral  acids  and  even  with  acetic  acid,  so 
that  it  cannot  be  precipitated  by  hydrogen  sulphide  (cf.  p.  254). 


490  COLLEGE    CHEMISTRY 

The  manganous  salts  of  weak  acids,  such  as  the  carbonate  and 
sulphide,  darken  when  exposed  to  air  and  are  oxidized,  with  forma- 
tion of  hydrated  manganic  oxide.  As  we  have  seen,  manganous 
hydroxide  is  similarly  oxidized  and  these  salts  are  precisely  the  ones 
which  should  furnish  the  hydroxide  by  hydrolysis.  While  there  is 
a  general  resemblance  between  the  manganous  salts  and  the  stannous. 
chromous,  and  ferrous  salts,  the  manganous  salts  of  active  acids  are 
not  oxidized  by  the  air  as  are  the  corresponding  salts  of  the  other 
three  metals. 

Manganic  Compounds.  —  The  base  of  this  set  of  compounds, 
manganic  hydroxide  Mn(OH)3,  is  slowly  deposited  by  the  action  of 
the  air  on  an  ammoniacal  solution  of  a  manganous  salt  in  salts  of 
ammonium.  Manganic  chloride  MnCl3  is  present  in  the  liquid  ob- 
tained by  the  action  of  hydrochloric  acid  upon  manganese  dioxide 
(cf.  p.  Ill),  but  loses  chlorine  very  readily. 

Manganites. —  Although  manganese  dioxide  interacts  when  fused 
with  potassium  hydroxide,  simple  salts  derived  from  H2MnO3(  =  H2O, 
MnO2)  or  H4MnO4(=  2H2O,MnO2)  are  not  formed.  The  products 
are  complex,  as  K2Mn5On.  Some  less  complex  manganites  are 
formed  in  the  Weldon  process  for  utilizing  the  manganous  chloride 
obtained  in  manufacturing  chlorine.  The  liquor  is  mixed  with 
slaked  lime,  and  air  is  blown  through  the  mass  of  calcium  and 
manganous  hydroxides  which  is  thus  obtained.  Black  manganites 
of  calcium,  such  as  CaMnO3(=  CaO,Mn02)  and  CaMn205(CaO, 
2MnO2)  are  thus  formed: 

Ca(OH)2  +  2Mn(OH)2  +  02-»CaMn2O5  4-  3H20, 

and  when  afterwards  treated  with  hydrochloric  acid  they  behave 
like  mixtures  of  manganese  dioxide  and  calcium  oxide. 

Manganates. —  When  one  of  the  oxides  of  manganese  is  fused 
with  potassium  carbonate  and  potassium  nitrate  a  green  mass  is 
obtained.  The  green  aqueous  extract  deposits  potassium  manganate 
K2MnO4  in  rhombic  crystals,  which  are  of  the  same  form  as  those  of 
potassium  sulphate,  and  are  almost  black: 

K2C03  +  Mn02  +  O  ->  K2MnO4  4-  CO2. 


MANGANESE  491 

The  acid  H2MnO4  is  itself  unknown.  The  potassium  salt  remains 
unchanged  in  solution  only  in  presence  of  free  alkali.  When  the 
concentration  of  the  hydroxide-ion  is  reduced  by  dilution,  or,  better 
still,  when  a  weak  acid  such  as  carbonic  acid  or  acetic  acid  is  used  to 
neutralize  it,  the  salt  is  decomposed  according  to  the  following 
equation: 

3K2MnO4  +  2H2O->4KOH  +  2KMnO4  +  MnO2. 

That  is,  a  precipitate  of  manganese  dioxide,  and  a  solution  of 
potassium  permanganate  are  obtained.  In  terms  of  the  ions  the 
equation  is  simpler: 

3Mn04"  +  2H'  -»  20H'  +  2MnO/  +  MnO2. 

Permanganates.  —  Potassium  permanganate  KMn04  is  made  by 
decomposition  of  the  manganate  as  shown  above,  and  is  obtained,  in 
purple  crystals  with  a  greenish  luster,  by  evaporation  of  the  solution. 
To  avoid  the  loss  of  manganese  thrown  down  as  dioxide,  the  action 
is  carried  out  commercially  by  passing  ozone  through  the  solution  of 
the  manganate:  2K2MnO4  +  O3  +  H2O  ->  2KMn04  +  02  +  2KOH. 
Sodium  permanganate  NaMnO4  is  made  in  a  similar  manner.  It  is  not 
obtainable  in  solid  form,  but  its  solution  is  known  as  "  Condy's  dis- 
infecting fluid."  This  liquid  owes  its  properties  to  the  oxidizing 
power  of  the  salt.  Permanganic  acid  is  a  very  active  acid,  that  is, 
it  is  highly  ionized  in  aqueous  solution.  A  solid  hydrate  of  the  acid 
may  be  secured  in  reddish-brown  crystals  by  adding  sulphuric  acid 
to  a  solution  of  barium  permanganate  and  allowing  the  filtrate  to 
evaporate: 

Ba(MnO4)2  +  H2SO4  +  zH2O  <=±  BaSO4  \  +  2HMnO4,zH2O. 

This  hydrate  decomposes,  on  being  warmed  to  32°,  and  yields  oxygen 
and  manganese  dioxide.  When  a  very  little  dry,  powdered  potas- 
sium permanganate  is  moistened  with  concentrated  sulphuric  acid, 
brownish-green,  oily  drops  of  permanganic  anhydride  (manganese  hept- 
oxide)  Mn2O7  are  formed.  This  compound  is  volatile,  giving  a  violet 
vapor,  and  is  apt  to  decompose  explosively  into  oxygen  and  manga- 
nese dioxide.  Its  oxidizing  power  is  such  that  combustibles 
like  paper,  ether,  and  illuminating-gas  are  set  on  fire  by  contact 
with  it. 


492  COLLEGE   CHEMISTRY 

Potassium  Permanganate  as  an  Oxidizing  Agent.  —  The 

actions  are  different  according  as  the  substance  is  employed  (1)  in 
acid,  or  (2)  in  neutral  solution. 

1.  In  presence  of  an  acid,  and  an  oxidizable  body,  a  manganous 
salt  is  always  formed.    The  schematic  equation,  Mn2O7  — >  2MnO  +  5O, 
shows  that  every  two  molecules  of  the  permanganate  yield  5O  for 
oxidizing  purposes.    Thus,  when  sulphuric  acid  is  added  to  potas- 
sium permanganate  solution,  and  sulphur  dioxide  is  led  through  the 
mixture,  we  have : 

2KMn04  +  3H,S04-+K,S04  +  2MnS04  +  3H,0(-f  50)  (1) 

(50)  +  5H2SO3  ->  5H2SO4 [ (2) 

2KMnO4+3H2SO4+5H2SO3->K2SO4  +  2MnS04  +  3H2O  +5H2SO4 

In  this  case,  since  sulphuric  acid  is  a  product,  the  preliminary  addi- 
tion of  the  acid  was  superfluous.  In  other  cases,  the  partial  equation 
(1),  showing  the  available  5O,  remains  the  same,  while  the  other  par- 
tial equation  varies  with  the  substance  being  oxidized.  Thus,  with 
hydrogen  sulphide  as  reducing  agent,  we  have: 

(0)  +  H2S  -»  H20  +  S     X  6  (20 

and  with  ferrous  sulphate,  we  get  ferric  sulphate: 

2FeS04  +  H2S04(+  O)  -»  Fe2(SO4)3  +  H2O     X  5         (2") 

As  before  (2')  and  (2")  must  be  multiplied  throughout  by  five,  before 
summation  is  made. 

2.  When  dry  potassium  permanganate  is  heated,  it  decomposes 

2KMn04  ->  K2Mn04  +  Mn02  +  02. 

The  neutral  solution  oxidizes  substances  which  are  reducing  agents. 
The  fingers  are  stained  brown  by  an  aqueous  solution,  receiving  a 
deposit  of  manganese  dioxide,  in  consequence  of  the  reducing  power 
of  the  unstable  organic  substances  in  the  skin.  The  destruction  of 
minute  organisms  by  Condy's  fluid  results  from  a  similar  action. 
When  the  powdered  salt  is  moistened  with  glycerine,  the  mass 
presently  bursts  into  flame. 

Analytical  Reactions  of  Manganese  Compounds  —  The  ions 
commonly  encountered  are  manganous-ion  Mn",  which  is  very  pare- 
pink  in  color,  permanganate-ion  MnO/,  which  is  purple,  and  man- 


MANGANESE  493 

ganate-ion  Mn04",  which  is  green.  The  manganous  compounds 
give  with  ammonium  sulphide  the  flesh-colored,  hydrated  sulphide 
which  is  soluble  in  acids.  Bases  give  the  white  hydroxide,  which 
darkens  by  oxidation,  and  is  soluble  in  salts  of  ammonium.  All 
compounds  of  manganese  confer  upon  the  borax  bead  an  amethyst 
color  which,  in  the  reducing  flame,  disappears.  A  bead  of  sodium 
carbonate  and  niter  becomes  green  on  account  of  the  formation  of 
the  manganate. 

Exercises.  —  1.  What  do  we  mean  by  saying  that,  (a)  chromous 
chloride  is  stable  (p.  81),  but  easily  oxidized  by  the  air,  (b)  perman- 
ganic acid  is  an  active  oxidizing  agent  in  presence  of  an  acid  (p.  194). 

2.  Formulate  the  oxidations  of  hydrogen  sulphide,  of  ferrous 
sulphate,  of  oxalic  acid  (to  carbon  dioxide),  and  of  nitrous  acicl  (to 
nitric  acid)  by  potassium  permanganate  in  acid  solution.  * 


CHAPTER  XLIII 
IRON,    COBALT,   NICKEL 

THE  elements  iron  (Fe,  at.  wt.  55.9),  cobalt  (Co,  at.  wt.  59),  and 
nickel  (Ni,  at.  wt.  58.7)  are  not  corresponding  members  of  successive 
periods,  like  the  families  hitherto  considered.  They  are  neighboring 
members  of  the  first  long  period,  lying  between  its  first  and  second 
octaves. 

IRON  Fe. 

Chemical  Relations  of  the  Element.  —  The  oxides  and 
hydroxides  FeO  and  Fe(OH)2,  Fe203  and  Fe(OH)3  are  basic,  the 
former  more  strongly  so  than  the  latter.  The  ferrous  salts,  derived 
from  Fe(OH)2,  resemble  those  of  the  magnesium  group  and  those  of 
Cr"  and  Mn**,  and  are  little  hydrolyzed.  The  ferric  salts,  derived 
from  Fe(OH)3,  resemble  those  of  Cr""  and  Al'"  and  are  hydrolyzed 
to  a  considerable  extent.  Iron  gives  also  a  few  ferrates  K2Fe04, 
CaFeO4,  etc.,  derived  from  an  acid  H2FeO4  which,  like  manganic 
acid  H2Mn04  (p.  491),  is  too  unstable  to  be  isolated.  Complex 
anions  containing  this  element,  such  as  the  anion  of  K4.Fe(CN)6,  are 
familiar,  but  complex  cations  containing  ammonia  are  unknown. 

Occurrence.  —  Free  iron  is  found  in  minute  particles  in  some 
basalts,  and  many  meteorites  are  composed  of  it.  Meteoric  iron 
can  be  distinguished  from  specimens  of  terrestrial  origin  by  the  fact 
that  it  contains  3-8  per  cent  of  nickel.  The  chief  ores  of  iron  are  the 
oxides,  hsematite  Fe203  and  magnetite  Fe304,  and  the  carbonate 
FeC03,  siderite.  The  first  is  reddish  and  radiated  in  structure;  but 
black,  shining,  rhombohedral  crystals,  known  as  specularite,  are  also 
found.  Hydrated  forms,  like  brown  iron  ore  2Fe2O3,  3H2O,  are  also 
common.  Siderite  is  pale-brown  in  color  and  rhombohedral,  like 
calcite.  When  mixed  with  clay  it  forms  iron-stone.  Pyrite  FeS2 
consists  of  golden-yellow,  shining  cubes  or  pentagonal  dodecahedra. 
It  is  used,  on  account  of  its  sulphur,  in  the  manufacture  of  sulphuric 

494 


IRON 


495 


acid,  but,  from  the  oxidized  residue,  iron  of  sufficient  purity  is 
obtained  with  difficulty.  Compounds  of  iron  are  contained  in 
chlorophyll  and  in  the  blood  (haemoglobin),  and  doubtless  play  an 
important  part  in  connection  with  the  vital  functions  of  these  sub- 
stances. Ammonium  sulphide  blackens  the  skin,  forming  ferrous 
sulphide  by  interaction  with  organic  compounds  of  iron  present  in 
the  tissues. 

Metallurgy.  —  The  ores  of  iron  are  first  roasted  in  order  to  decom- 
pose carbonates  and  oxidize  sulphides,  if  these  salts  are  present.  Coke 
is  then  used  to  reduce  the  oxides. 
Ores  containing  lime  or  magnesia  are 
mixed  with  an  acid  flux,  such  as  sand 
or  clay-slate,  in  order  that  a  fusible 
slag  may  be  formed.  Conversely,  ores 
containing  silica  and  clay  are  mixed 
with  limestone.  With  proper  adjust- 
ment of  the  ingredients  the  process  can 
be  carried  on  continuously  in  a  blast 
furnace  (Fig.  66).  The  solid  materials 
thrown  in  at  the  top  are  converted,  as 
they  slowly  descend,  completely  into 
gases  which  escape  and  liquids  (iron  and 
slag)  which  are  tapped  off  at  the 
bottom.  Heated  air  is  blown  in  at  the 
bottom  through  tuyeres,  and  the  top  is 
closed  by  a  bell  which  descends  for  a 
moment  when  an  addition  is  made  to 
the  charge.  The  gases,  which  contain 
much  carbon  monoxide,  are  led  off  and 
used  to  heat  the  blast  or  to  drive  gas-  Y\G.  66. 

engines. 

The   main   action  takes   place   between   the   carbon   monoxide, 
present  in  consequence  of  the  excess  of  carbon,  and  the  oxide  of  iron : 

Fe3O4  +  4CO  «=±  3Fe  +  4CO2. 

Since  the  action  is  a  reversible  one,  a  large  excess  of  carbon  monoxide 
is  required. 

In  the  upper  part  of  the  furnace,  the  heat  (400°)  loosens  the 
texture  of  the  ore.     Further  down,  the  temperature  is  higher  (500- 


496  COLLEGE   CHEMISTRY 

900°),  and  the  carbon  monoxide  reduces  the  oxide  of  iron  to  particles 
of  soft  iron.  A  temperature  high  enough  to  melt  pure  iron  is  barely 
reached  anywhere  in  the  furnace,  but,  a  little  lower  down,  by  union 
with  carbon,  the  more  fusible  cast  iron  (1200°)  is  formed  and  falls  in 
drops  to  the  bottom.  It  is  in  this  region  also  that  the  slag  is  pro- 
duced. If  the  flux  had  begun  sooner  to  interact  with  the  unreduced 
ore,  iron  would  have  been  lost  by  the  formation  of  the  silicate.  The 
iron  collects  below  the  slag,  and  the  latter  flows  continuously  from 
a  small  hole.  The  former  is  tapped  off  at  intervals  of  six  hours  or  so 
from  a  lower  opening. 

Pure  iron  may  be  made  by  reducing  the  purified  oxalate  in  a  stream 
of  hydrogen,  or  by  Goldschmidt's  method  (p.  444). 

Cast  Iron  and  Wrought  Iron.  —  Pure  iron  is  not  manu- 
factured, and  indeed  would  be  too  soft  for  most  purposes.  Piano- 
wire,  however,  is  about  99.7  per  cent  pure.  The  product  obtained 
from  the  blast  furnace  contains  92-94  per  cent  of  iron  along  with 
2.6-4.3  per  cent  of  carbon,  often  nearly  as  much  silicon,  varying 
proportions  of  manganese,  and  some  phosphorus  and  sulphur.  The 
last  four  ingredients  are  liberated  from  combination  with  oxygen  by 
the  carbon  in  the  hottest  part  of  the  furnace  and  combine  or  alloy 
themselves  with  the  iron.  Cast  iron  does  not  soften  before  melting, 
as  does  the  purer  wrought  iron,  but  melts  sharply  at  1150-1250° 
according  to  the  amount  of  foreign  material  it  contains.  When 
suddenly  cooled  it  gives  chilled  cast  iron  which  is  very  brittle  and 
looks  homogeneous  to  the  eye,  all  the  carbon  being  present  in  the 
form  of  carbide  of  iron  Fe3C  in  solid  solution  in  the  metal.  This 
solid  solution  is  exceedingly  hard.  By  slower  cooling,  time  is  per- 
mitted for  the  separation  of  part  of  the  carbon  as  graphite,  which 
appears  in  tiny  black  scales,  and  gray  cast  iron  results.  This  mixture 
is  much  softer,  on  account  of  the  amount  of  free,  relatively  pure  iron 
which  it  contains.  Spiegel  iron  is  cast  iron  made  from  ores  contain- 
ing 5-20  per  cent  of  manganese  and  the  usual  proportion  of  carbon. 

Wrought  iron  is  made  by  heating  the  broken  "  pigs  "  of  cast  iron 
upon  a  layer  of  material  containing  oxide  of  iron  and  hammer-slag 
(basic  silicate  of  iron)  spread  on  the  bed  of  a  reverberatory  furnace. 
The  carbon,  silicon,  and  phosphorus  combine  with  the  oxygen  of  the 
oxide,  and  the  last  two  pass  into  the  slag.  The  sulphur  is  found  in 
the  slag  as  ferrous  sulphide.  On  account  of  the  effervescence  due 


IRON  497 

to  the  escape  of  carbon  monoxide,  the  process  is  called  "  pig-boiling." 
The  iron  is  stirred  with  iron  rods  ("  puddled  ")  and  stiffens  as  it 
becomes  purer,  until  finally  it  can  be  withdrawn  in  balls  ("  blooms  ") 
and  freed  from  slag  under  the  steam-hammer  or  by  rolling.  It  now 
softens  sufficiently  for  welding  below  1000°  and  melts  at  1550°. 
Wrought  iron  should  not  contain  more  than  0.15  per  cent  of  carbon. 
The  above  operations  are  now  largely  performed  by  machinery. 

Steel.  —  This  is  a  variety  of  iron  almost  free  from  phosphorus, 
sulphur,  and  silicon.  Tool-steel  contains  0.9-1.5  per  cent  of  carbon, 
structural  steel  only  0.2-0.6  per  cent,  and  mild  steel  0.2  per  cent  or 
even  less.  Steel  combines  the  properties  of  cast  and  of  wrought  iron, 
being  hard  and  elastic,  and  at  the 
same  time  available  for  forging 
and  welding  when  the  proportion 
of  carbon  is  not  too  high. 

Steel  is  made  largely  by  the 
Bessemer  process.  The  molten 
cast  iron  is  poured  into  a  con- 
verter (Fig.  67)  and  a  blast  of  air 
(a)  is  blown  through  it.  The  oxi-  PIG 

dation  of  the  manganese,  carbon, 

silicon,  and  a  little  of  the  iron  gives  out  sufficient  heat  to  raise  the 
temperature  of  the  mass  above  the  melting-point  of  wrought  iron. 
The  required  proportion  of  carbon  is  then  introduced  by  adding  pure 
cast  iron,  spiegel  iron,  or  coke,  and  the  contents,  first  the  slag,  and 
then  the  molten  steel,  are  finally  poured  out  by  turning  the  con- 
verter. When  the  cast  iron  contains  much  phosphorus,  the  oxide 
of  this  element  is  reduced  again  by  the  iron  as  fast  as  it  is  formed  by 
the  blast.  In  such  cases  a  basic  lining  containing  lime  and  magnesia 
takes  the  place  of  the  sand  and  clay  lining  of  the  ordinary  Bessemer 
converter,  and  a  slag  containing  a  basic  phosphate  of  calcium  is  pro- 
duced. This  modification  constitutes  what  is  known  as  the  Thomas- 
Gilchrist  process.  The  slag  ("  Thomas-slag ")  when  pulverized 
forms  a  valuable  fertilizer  (cf.  p.  401). 

In  the  Siemens-Martin,  or  open  hearth  process,  the  cast  iron  is 
melted  in  a  saucer-shaped  depression  lined  with  sand,  and  scraps  of 
iron  plate  (for  dilution)  and  haematite,  or  some  other  oxide  ore,  are 
then  added  in  proper  proportions.  The  materials  are  heated  with 


498  COLLEGE    CHEMISTRY 

gas  fuel  for  8-10  hours  until  a  sample  shows,  under  the  hammer,  that 
the  process  is  complete.  The  product  is  then  drawn  off  through  a 
hole  and  cast  in  molds. 

When  steel  is  heated  to  redness  and  cooled  slowly,  it  is  compara- 
tively soft.  Sudden  chilling,  however,  renders  it  harder  than  glass. 
The  explanation  of  this  behavior  is  the  same  as  is  the  case  of  cast 
iron  (p.  496).  By  subsequent,  cautious  heating  the  hardness  may 
be  reduced  to  any  required  extent,  and  this  treatment  is  called 
"  tempering." 

Chemical  Properties. — When  exposed  to  moist  air,  iron  receives 
a  loosely  adherent  coating  of  rust  (2Fe2O3,Fe(OH)3).  This  product 
may  be  formed  by  displacement  of  the  hydrogen  of  carbonic 
acid,  the  oxygen  assisting  (cf.  p.  415),  and  subsequent  hydrolysis  of 
the  carbonate  and  oxidation  of  the  ferrous  hydroxide.  The  fact 
that  alkalies  prevent  rusting  favors  this  view,  for  they  should 
diminish  the  amount  of  hydrogen-ion.  Many  chemists  believe, 
however,  that  carbonic  acid  plays  no  part  in  the  process. 

Iron  burns  in  oxygen  and  it  interacts  with  superheated  steam, 
in  both  cases  giving  Fe304.  A  superficial  layer  of  this  oxide  ad- 
heres firmly  and  protects  the  iron  from  the  action  of  the  air  (Barff  s 
process  for  prevention  of  rusting). 

Iron  displaces  hydrogen  easily  from  dilute  acids.  Steel  and  cast 
iron,  which  contain  iron,  its  carbide,  and  graphite,  give  with  cold 
dilute  acids  almost  pure  hydrogen,  and  the  carbide  and  graphite 
remain  unattacked.  More  concentrated  acids,  however,  particularly 
when  warm,  generate,  along  with  hydrogen,  hydrocarbons  formed  by 
interaction  with  the  carbide  (p.  330).  The  odor  of  the  gas  is  due  to 
compounds  of  sulphur  and  phosphorus. 

Ferrous  Compounds. — Ferrous  chloride  is  obtained  as  a  pale- 
green  hydrate  FeCl2,4H2O  by  interaction  of  hydrochloric  acid  with 
the  metal  or  the  carbonate.  The  anhydrous  salt  sublimes  in  color- 
less crystals  when  hydrogen  chloride  is  led  over  the  heated  metal. 
In  solution  the  salt  is  oxidized  by  the  air  to  a  basic  ferric  chloride: 

4Fe"  +  02  +  2H20  -»  4Fe"'  +  4OH'. 

In  presence  of  excess  of  the  acid,  normal  ferric  chloride  is  formed. 
With  nitric  acid,  ferric  chloride  and  nitric  oxide  are  produced  (p.  295). 


IRON  499 

Ferrous  hydroxide  Fe(OH)2  is  thrown  down  as  a  white  precipitate, 
but  rapidly  becomes  dirty-green  and  finally  brown,  by  oxidation.  It 
dissolves  in  solutions  of  salts  of  ammonium,  being,  like  magnesium 
hydroxide  (p.  429) ,  sufficiently  soluble  in  water  to  require  an  appre- 
ciable concentration  of  OH7  for  its  precipitation.  The  ammonium 
salts  convert  this  into  molecular  ammonium  hydroxide.  Ferrous 
oxide  FeO  is  black,  and  is  formed  by  heating  ferrous  oxalate  in 
absence  of  air.  It  may  be  made  also  by  cautious  reduction  of  ferric 
oxide  by  hydrogen  (at  about  300°),  but  is  easily  reduced  further 
to  the  metal.  It  catches  fire  spontaneously  wfeten  exposed  to  the 
air. 

Ferrous  carbonate  FeCO3  is  found  in  nature,  and  may  be  made  in 
slightly  hydrolyzed  form  by  precipitation.  The  precipitate  is  white 
but  rapidly  darkens  and  finally  becomes  brown,  the  ferrous  hydrox- 
ide produced  by  hydrolysis  being  oxidized  to  the  ferric  condition. 
The  salt  interacts  with  water  containing  carbonic  acid,  after  the 
manner  of  calcium  carbonate  (p.  324),  and  hence  is  found  in  solution 
in  natural  (chalybeate)  waters. 

Ferrous  sulphide  FeS  may  be  formed  as  a  black,  metallic-looking 
mass  by  heating  together  the  free  elements.  It  is  produced  by 
precipitation  with  ammonium  sulphide,  but  not  with  hydrogen  sul- 
phide. It  interacts  readily  with  dilute  acids.  The  precipitated 
form  is  slowly  oxidized  to  ferrous  sulphate  by  the  air. 

Ferrous  sulphate  is  obtained  by  allowing  pyrites  to  oxidize  in  the 
air  and  leaching  the  residue: 

2FeS2  +  702  +  2H2O  ->  2FeS04  +  2H2SO4. 

The  liquor  is  treated  with  scrap  iron  and  the  neutral  solution  evapo- 
rated until  a  hydrate  FeS04,7H2O,  green  vitriol,  or  "  copperas,"  is 
deposited.  The  crystals  are  efflorescent,  and  become  also  brown 
from  oxidation  to  a  basic  ferric  sulphate: 

4FeS04  +  02  +  2H2O  -»  4Fe(OH)SO4. 

With  excess  of  sulphuric  acid  and  air,  or  an  oxidizing  agent,  such  as 
nitric  acid,  ferric  sulphate  is  formed.  The  ferrous  sulphate  is  used 
in  dyeing  and  in  making  writing-ink.  The  extract  of  nut-galls  con- 
tains tannic  acid,  HC14H9Ott,  which,  with  ferrous  sulphate,  gives 
ferrous  tannate.  A  solution  of  this  salt  containing  gum-arabic 
and  some  dark  coloring  matter  constitutes  the  ink.  When  the 


TP* 

^500  COLLEGE   CHEMISTRY 

writing  is  exposed  to  the  air,  the  ferrous  tannate  is  oxidized  to 
the  ferric  condition,  and  the  ferric  compound  is  a  fine,  black 
precipitate  (cf.  p.  420). 

Ferric  Compounds.  —  By  leading  chlorine  into  a  solution  of 
ferrous  chloride,  and  evaporating  until  the  proper  proportion  of 
water  alone  remains,  a  yellow,  deliquescent  hydrate  of  ferric  chloride, 
FeCl3,  6H2O  is  obtained.  When  this  is  heated  still  further,  hydroly- 
sis takes  place  and  the  oxide  remains.  When  chlorine  is  passed  over 
heated  iron,  the  anhydrous  salt  sublimes  in  dark  scales,  which  are 
red  by  transmitted  light.  In  solution,  the  salt,  like  other  ferric  salts, 
can  be  reduced  to  the  ferrous  condition  by  boiling  with  iron.  The 
same  reduction  is  effected  by  hydrogen  sulphide: 

2Fe"*  +  Fe— »3Fe". 
2Fe'"  +  S"->2Fe"  +  S  J. 

The  ferric  ioi^is  almost  colorless,  the  yellow-brown  color  of  solutions 
of  ferric  chloride  being  due  to  the  presence  of  ferric  hydroxide  pro- 
duced by  hydrolysis.  The  color  deepens  when  the  solution  is  heated, 
and  fades  again  very  slowly,  by  reversal  of  the  action,  when  the  cold 
solution  is  allowed  to  stand. 

Ferric  hydroxide  Fe(OH)3  appears  as  a  brown  precipitate  when  abase 
is  added  to  a  ferric  salt.  It  does  not  interact  with  excess  of  the  alkali. 
In  this  form  the  substance  dries  to  the  oxide  without  giving  definite 
intermediate  hydrated  oxides.  The  hydrates,  Fe2O3,2Fe(OH)3 
(brown  iron  ore)  and  Fe2O3,4Fe(OH)3  (bog  iron  ore),  however,  are 
found  in  nature  (see  Rust,  p.  498).  The  hydroxide  passes  easily 
into  colloidal  solution  in  a  solution  of  ferric  chloride,  and  by  subse- 
quent dialysis  through  a  piece  of  parchment  the  salt  can  be  separated, 
and  a  pure  aqueous  solution  of  the  hydroxide  obtained.  This  solution 
is  red  in  color,  shows  no  depression  in  the  freezing-point,  and  is  not 
an  electrolyte.  It  deposits  the  hydroxide  as  a  brown  precipitate 
when  various  substances  are  added  to  the  solution. 

Ferric  oxide,  Fe203,  is  sold  as  "  rouge  "  and  "  Venetian  red."  It  is 
made  from  the  ferrous  sulphate,  obtained  in  cleaning  iron  ware  which 
is  to  be  tinned  or  galvanized,  and  in  other  ways  in  the  arts.  The 
salt  is  allowed  to  oxidize,  and  the  ferric  hydroxide,  thrown  down  by 
the  addition  of  lime,  is  calcined.  This  oxide  is  not  distinctly  acidic, 
but  by  fusion  with  more  basic  oxides,  compounds  like  franklinite 


IRON  501 

Zn(FeO2)2  may  be  formed.  It  is  reduced  by  hydrogen,  at  about 
300°  to  ferrous  oxide  (p.  499),  and  at  700-800°  to  metallic  iroru 

Magnetic  oxide  of  iron  Fe3O4  or  lodestone  is  found  in  nature,  and 
is  formed  by  the  action  of  air  (hammer-scale),  steam,  or  carbon 
dioxide  on  iron.  It  forms  octahedral  crystals. 

Ferric  sulphide  Fe2S3  may  be  made  by  fusing  together  the  free 
elements,  but  is  not  obtained  by  precipitation.  Soluble  sulphides 
first  reduce  the  ferric  salt  to  the  ferrous  conditions,  liberating  sulphur, 
Fe2(S04)3  +  (NH4)2S  ->  2FeSO4  +  (NH4)2S04  +  S, 

and  then  give  ferrous  sulphide  (p.  254). 

Ferric  sulphate  Fe2(S04)3  is  formed  by  oxidation  of  ferrous  sulphate, 
and  is  obtained  as  a  white  mass  by  evaporation.  It  gives  alums, 
such  as  (NH4)2SO4,Fe2(SO4)3,24H2O,  which  are  almost  colorless  when 
pure,  but  usually  have  a  pale  reddish-violet  tinge. 


Pyrite.  —  The  mineral  pyrite  FeS2  (Fools'  gold)  is  the  sulphide  of 
iron  which  is  most  stable  in  the  air.  It  is  found  in  nat^^  in  the  form 
of  glittering,  golden-yellow  cubes,  octahedrons,  \\M  pentagonal 
dodecahedrons.  It  is  not  attacked  by  dilute  &c\f  but  concen- 
trated hydrochloric  acid  slowly  converts  it  into  ferrous  chloride  and 
sulphur.  It  is  reduced  by  hydrogen  to  ferrous  sulphide. 


Cyanides* —  When  potassium  cyanide  is  added  to  solutions  of 
ferrous  or  ferric  salts,  yellowish  precipitates  are  produced,  but  the 
simple  cyanides  cannot  be  obtained  in  pure  form.  These  precipi- 
tates interact  with  "excess  of  the  cyanide  giving  soluble  complex 
cyanides  of  the  forms  4KCN,Fe(CN)2  and  3KCN,Fe(CN)3.  These 
are  called  ferro-  and  ferricyanide  of  potassium,  respectively. 

Ferrocyanide  of  potassium  K4Fe(CN)6,  3H20,  "  yellow  prussiate  of 
potash,"  is  made  by  heating  nitrogenous  animal  refuse,  such  as  blood, 
with  iron  filings  and  potassium  carbonate.  The  resulting  mass 
contains  potassium  cyanide  and  ferrous  sulphide,  and  when  it  is 
treated  with  warm  water  these  interact  and  produce  the  ferrocyanide: 

2KCN  +  FeS  ->  Fe(CN)2  +  K2S, 
4KCN  +  Fe(CN),-»K4.Fe(CN)6. 

The  salt  is  made  also  from  the  cyanogen  contained  in  crude  illumi- 
nating-gas. The  trihydrate  forms  large,  yellow,  monosymmetric 
tables.  The  solution  contains  almost  exclusively  the  ions  K"  and 


502  COLLEGE    CHEMISTRY 

Fe(CN)6"",  and  gives  none  of  the  reactions  of  the  ferrous  ion  Fe". 
The  corresponding  acid  H4.Fe(CN)6  may  be  obtained  as  white 
crystalline  scales  by  addition  of  an  acid  and  of  ether  (in  which  the 
substance  is  less  soluble  than  in  water)  to  the  salt.  The  acid  is  a 
fairly  active  one,  but  is  unstable  and  decomposes  in  a  complex 
manner.  Other  ferrocyanides  m^-he.  made  by  precipitation.  That 
of  copper  Cu2.Fe(CN)$s  brown,  and  temfrferrocyanide  Fe4[Fe(CN)6]3 
has  a  brilliant  ,blue  color  ("  Prussian  blue  IjLThe  ferrous  com- 
pound Fe2Fe^€)N)6,  or  perhaps  K2FeFe(CN)6,  iP^ite  but  quickly 
becomes  blu£  by  oxidation.  The  ferrocyanides  are  not  poisonous. 

Ferricyajiide  of  potassium  K3Fe(CN)6  is  easily  made  from  the  ferro- 
cyanide  by  oxidation : 

£      2K4Fe(CN)6  +  C12  ->  2KC1  +  2K3.Fe(CN)tj 
or  2Fe(CN)6""  +  C12  -»  2Fe(CN)6///  +  2(1'. 

It  forms  red  monosymmetric  prisms.  The  free  acid  !Jfe(CN)6  is 
unstable.  Qjfcer  salts  may  be  prepared  by  precipitatijlp  Ferrous 
ferricyanide J^[Fe(CN)6]2  is  deep-blue  in  color  ("  Turnfull's  blue  ")• 
With  ferrie  s$ts  only  a  brown  solution  is  obtained,  jfc 


Iron  Cfirboif^fi.  —  When  carbon  monoxidjKs  led  over  finely 
divided  iron  atlBJk)0,  or-.^nder  eight  atmgsfftj^ps  pressure  at  the 
ordinary  temperamrej^iQi^J^l^^cjsmjpcRiiHls-  of  the  composition 
Fe(CO)4,  the  tetracarbonyl,  and  Fe(CO%,  the  pentacarbonyl,  are 
formed.  When  the  gaseous  mixture  is  heated^ more  strongly,  the 
compounds  decompose  again,  and  iron  is  deposited.  Illuminating- 
gas  burners  frequently  receive  a  deposit  of  iron  from  this  cause. 

Analytical  Reactions  of  Compounds  of  Iron.  —  There  are 
two  ionic  forms  of  iron,  ferrous-ion  Fe",  which  is  very  pale-green,  and 
ferric-ion  Fe'",  which  is  almost  colorless.  Ammonium  sulphide 
forms,  with  both,  black  ferrous  sulphide  which  is  soluble  in*  dilute 
acids.  The  hydroxides  are  white  and  brown  respectively,  and 
ferrous  carbonate  is  white.  With  ferric  salts,  soluble  carbonates 
yield  the  hydroxide.  With  ferrocyanide  of  potassium,  ferrous  salts 
give  a  white,  and  ferric  salts  a  blue,  precipitate.  With  ferricyanide 
of  potassium  the  former  give  a  deep-blue  precipitate,  and  the  latter 
a  brown  solution.  Ferric  thiocyanate  Fe(CNS)3  is  deep-red  (p.  IS9). 


COBALT  503 

With  borax,  iron  compounds  give  a  bead  which  is  green  in  the 
reducing  flame,  and  colorless  or,  with  much  iron,  yellow  or  even 
brown  when  oxidized. 

COBALT  Co. 

The  Chemical  Relations  of  the  Element.  —  Cobalt  forms 
cobaltous  and  cobaltic  oxides  and  hydroxides  CoO  and  Co(OH)2, 
Co2O3  and  Co(OH)3,  respectively,  which  are  all  basic,  the  former 
more  so  than  the  latter.  The  cobaltous  salts  are  little  hydrolyzed, 
but  the  cobaltic  salts  are  completely  decomposed  by  water.  The 
latter  also  liberate  readily  one-third  of  the  negative  radical,  after  the 
manner  of  manganic  salts,  becoming  cobaltous.  Complex  cations 
and  anions  containing  cobalt  are  very  numerous  and  very  stable. 

Occurrence  and  Properties. —  Cobalt  is  found  along  with  nickel 
in  smaltite  CoAs2  and  cobaltite  CoAsS.  The  pure  metal  may  be 
made  by  Goldschmidt's  process,  or  by  reducing  the  oxalate,  or  an 
oxide,  with  hydrogen. 

The  metal  is  silver-white,  with  a  faint  suggestion  of  pink.  It  is 
less  tough  thafi  iron,  and  has  no  commercial  applications.  It  dis- 
places hydrogen  slowly  from  dilute  acids,  but  interacts  readily  with 
nitric  acid. 

Cobaltous  Compounds.  —  The  chloride  CoCl2, 6H2O  may  be  made 
by  treating  the  oxide  with  hydrochloric  acid.  It  forms  red  prisms, 
and  when  partiallyjgibfijjiipletely  dehydrated  becomes  deep-blue. 
Writing  made  with  SoBBed  solution  upon  paper  is  almost  invisible, 
but  becomes  blue  when  warmed  and  afterwards  takes  up  moisture 
from  the  air,  aj^l  is  once  more  invisible  ("  sympathetic  ink  ").  Most 
cobaltous  compounds  are  red  when  hydrated  or  in  solution  (Co") 
and  blue  when  dehydrated.  By  addition  of  sodium  hydroxide  to  a 
cobaltous  salt,  a  blue  basic  salt  is  precipitated.  When  the  mixture 
is  boiled,  the  red  cobaltous  hydroxide,  Co(OH)2  is  formed.  This 
becomes  brown  through  oxidation  by  the  air.  It  interacts  with 
ammonium  hydroxide,  giving  a  soluble  ammonio-cobaltous  hydrox- 
ide, which  is  quickly  oxidized  by  the  air  to  an  ammonio-cobaltic 
compound  (see  below).  It  dissolves  also  in  salts  of  ammonium  as 
magnesium  hydroxide  does  (p.  429).  When  dehydrated  it  leaves 
the  black  cobaltous  oxide  CoO.  Cobaltous  sulphate,  CoSO4,  7H2O, 


604  COLLEGE   CHEMISTRY 

and  cobaltous  nitrate,  Co(N03)2,  6H2O,  are  familiar  salts.  The  black 
cobaltous  sulphide,  CoS,  is  precipitated  by  ammonium  sulphide  from 
solutions  of  all  salts,  and  even  by  hydrogen  sulphide  from  the  acetate, 
or  a  solution  containing  much  sodium  acetate  (cf.  p.  399).  Once  it 
has  been  formed,  it  does  not  interact  even  with  dilute  hydrochloric 
acid,  having  apparently  changed  into  a  less  active  form.  A  sort  of 
cobalt  glass,  made  by  fusing  sand,  cobalt  oxide,  and  potassium 
nitrate,  forms,  when  powdered,  a  blue  pigment  ("  smalt  ")  used  in 
china-painting  and  by  artists. 

Cobaltic  Compounds.  —  By  addition  of  a  hypochlorite  to  a  solu- 
tion of  a  cobaltous  salt,  cobaltic  hydroxide  Co(OH)3,  a  black  powder, 
is  precipitated.  Cautious  ignition  of  the  nitrate  gives  cobaltic  oxide, 
Co2O3.  Stronger  ignition  gives  the  commercial  oxide,  which  is  a 
cobalto-cobaltic  oxide  Co304.  Cobaltic  oxide  dissolves  in  cold  hydro- 
chloric acid,  but  the  solution  gives  off  chlorine  when  warmed.  By 
placing  cobaltous  sulphate  round  the  anode  of  an  electrolytic  cell, 
crystals  of  cobaltic  sulphate,  Co2(S04)3,  have  been  made  and  cobaltic 
alums  have  also  been  prepared. 

Complex  Compounds.  —  Potassium  cyanide  precipitates  from 
cobaltous  salts  a  brownish-white  cyanide  which  interacts  with  excess 
of  the  reagent,  giving  a  solution  of  potassium  cobaltocyanide  K4.Co 
(CN)6  (cf.  p.  501).  This  compound  is  easily  oxidized  by  chlorine,  or 
even  when  the  solution  is  boiled  in  the  air,  and  the  colorless  potassium 
cobalticyanide  is  formed: 

4K4Co(CN)6  +  2H20  +  O2->  4K3.Co(CN)6  +  4KOH. 

The  solution  gives  none  of  the  reactions  of  Co"*,  and  with  acids  the 
very  stable  cobalticyaiiic  acid,  H3Co(CN)3,  is  liberated. 

When  acetic  acid  and  potassium  nitrite  are  added  to  a  cobaltous 
salt,  the  latter  is  oxidized  by  the  nitrous  acid  (liberated  by  the  acetic 
acid)  and  a  white  complex  salt  K3.Co(NO2)6,nH2O  (==  Co(N02)3, 
3KNO2),  potassium  cobaltinitrite,  is  thrown  down. 

Cobaltic  salts  give  with  ammonia  complex  compounds  which  are 
many  and. various.  The  cations  often  contain  negative  groups,  and 
are  such  as  Co(NH3)6"',  Co(NH3)5Cl"  and  Co(NH3)5N02".  Usually 
the  solutions  give  none  of  the  reactions  of  cobaltic  ions,  and  often 
fail  likewise  to  give  those  of  the  anion  of  the  original  salt. 

H  * 


NICKEL  506 

NICKEL. 

The  Chemical  lielations  of  the  Element.  —  Nickel  forms 
nickelous  and  nickelic  oxides  and  hydroxides  NiO  and  Ni(OH)2, 
Ni2O3,  and  Ni(OH)3,  but  only  the  former  are  basic.  The  nickelous 
salts  resemble  the  cobaltous  and  ferrous  salts,  but  are  not  oxidizable 
into  corresponding  nickelic  compounds.  Since  there  are  no  nickelic 
salts,  there  are  here  no  analogues  of  the  cobalticyanides  or  the 
cobaltinitrites.  The  complex  nickelous  salts,  like  the  complex 
cobaltous  salts,  and  unlike  the  complex  cobaltic  salts,  are  unstable, 
and  so  give  some  of  the  reactions  of  Ni". 

Occurrence  and  Properties.  —  Nickel  occurs  free  in  meteorites 
and  in  niccolite  NiAs  and  nickel  glance  NiAsS.  It  is  now  manu- 
factured chiefly  from  pentlandite  [Ni,Fe]S  and  other  minerals  found 
at  Sudbury  (Ontario),  and  from  garnierite,  a  silicate  of  nickel  and 
magnesium  found  in  New  Caledonia. 

The  metal  is  white,  with  a  faint  tinge  of  yellow,  is  very  hard,  and 
takes  a  high  polish.  It  is  used  in  making  alloys,  such  as  German 
silver  (copper,  zinc,  nickel,  2:1:1)  and  the  "  nickel  "  used  in  coin- 
age (copper,  nickel,  3:1).  Nickel  plating  on  iron  is  accomplished 
exactly  like  silver  plating  (p.  421).  The  bath  contains  an  ammonia- 
cal  solution  of  ammonium-nickel  sulphate  (NH4)2SO4,NiS04,  6H20, 
and  a  plate  of  nickel  forms  the  anode. 

The  metal  rusts  very  slowly  in  moist  air.  It  displaces  hydrogen 
with  difficulty  from  dilute  acids  but  interacts  with  nitric  acid. 

Compounds  of  Nickel.  —  The  chloride  NiCl2,6H20,  is  made  by 
treating  any  of  the  oxides  with  hydrochloric  acid,  and  is  green  in 
color  (when  anhydrous,  brown).  The  sulphate,  NiSO4,6H20,  which 
crystallizes  in  square  prismatic  forms  at  30-40°,  is  the  most  familiar 
salt.  Nickelous  hydroxide,  Ni(OH)2,  is  formed  as  an  apple-green 
precipitate,  and  when  heated  leaves  the  green  nickelous  oxide,  NiO. 
It  dissolves  in  ammonium  hydroxide,  giving  a  complex  nickel- 
ammonia  cation.  It  is  soluble  also  in  salts  of  ammonium  (cf.  p.  429). 
By  cautious  ignition  of  the  nitrate,  nickelic  oxide,  Ni203,  is  formed  as 
a  black  powder.  The  oxides  and  salts,  when  heated  strongly  in 
oxygen  give  the  oxide  Ni3O4.  The  last  two  oxides  liberate  chlorine 
when  treated  with  hydrochloric  acid,  and  give  nickelous  chloride. 
Nickelic  hydroxide,  Ni(OH)3,  is  a  black  precipitate  formed  when  a 


506  COLLEGE  CHEMISTRY 

hypochlorite  is  added  to  any  salt  of  nickel.  Nickelous  sulphide  is 
thrown  down  by  ammonium  sulphide,  and  behaves  like  cobaltous 
sulphide  (p.  504).  It  forms  a  brown  colloidal  solution  when  excess 
of  the  precipitant  is  used,  and  is  then  deposited  very  slowly. 

With  potassium  cyanide  and  a  salt  of  nickel  the  greenish  nickelous 
cyanide,  Ni(CN)2,  is  first  precipitated.  This  dissolves  in  excess  of 
the  reagent,  and  a  complex  salt  K2Ni(CN)4,  H2O(=  2KCN,Ni(CN)2) 
may  be  obtained  from  the  solution.  This  salt  is  of  different  compo- 
sition from  the  corresponding  compounds  of  cobalt  and  of  iron,  and 
is  less  stable.  Thus,  with  bleaching  powder,  it  gives  Ni(OH)3  as  a 
black  precipitate.  When  the  solution  is  boiled  in  the  air  no  oxida- 
tion to  a  complex  nickelicyanide  occurs,  and  indeed  no  such  salts  are 
known.  This  fact  enables  the  chemist  to  separate  cobalt  and  nickel, 
for  when  the  mixed  cyanides  are  boiled  and  then  treated  with  bleach- 
ing powder,  the  cobalticyanide  is  unaffected.  With  potassium 
nitrite  and  acetic  acid  no  insoluble  compound  corresponding  to  that 
given  by  cobalt  salts  is  formed  by  salts  of  nickel.  The  only  known 
compound  which  could  be  formed,  4KN02,Ni(N02)2,  is  soluble. 
This  action  also  is  used  for  the  purpose  of  separation. 

When  finely  divided  nickel,  made  by  reducing  the  oxide  or  oxalate 
with  hydrogen  at  a  moderate  temperature,  is  exposed  to  a  stream  of 
cold  carbon  monoxide,  nickel  carbonyl  Ni(CO)4  is  formed.  This  is  a 
vapor  and  is  condensable  to  a  colorless  liquid  (b.-p.  43°  and  m.-p. 
-  25°).  The  vapor  is  poisonous.  When  heated  to  150-180°  it  is 
dissociated  and  nickel  is  deposited.  Cobalt  forms  no  corresponding 
compound. 

Analytical  Reactions  of  Compounds  of  Cobalt  and  Nickel. 

— The  cobalt  ion  Co"  is  pink,  and  the  nickelous  ion  Ni"  green.  The 
reactions  used  in  analysis  have  been  described  in  the  preceding 
paragraphs.  With  borax,  cobalt  compounds  give  a  blue  bead,  and 
nickel  compounds  a  bead  which  is  brown  in  the  oxidizing  flame  and 
cloudy,  from  the  presence  of  gray,  metallic  nickel,  when  reduced. 

Exercises. —  1.  What  would  be  the  interactions  of  calcium  car- 
bonate when  fused  with  sand  and  with  clay,  respectively? 

2.  Make  equations  representing,  (a)  the  oxidation  of  ferrous 
chloride  by  air,  (6)  the  hydrolysis  of  ferrous  carbonate  and  the 
oxidation  of  ferrous  hydroxide,  (c)  the  oxidation  of  ferrous  sulphate 


NICKEL  507 

with  excess  of  sulphuric  acid  by  hypochlorous  acid,  (d)  the  formation 
of  ferrous  and  ferric  tannates  (p.  499),  (e)  the  reduction  of  ferric 
chloride  by  iron  and  by  hydrogen  sulphide,  respectively,  (/)  the  dry 
distillation  of  basic  ferric  sulphate,  (g)  the  formation  of  ferric  ferro- 
cyanide  and  of  ferrous  ferricyanide. 

3.  Explain  the  solubility  of  cobaltous  and  nickelous  hydroxide  in 
salts  of  ammonium. 

4.  Construct  equations  to  show  the  formation,  (a)  of  the  insoluble 
potassium  cobalti nitrite  (nitric  oxide  is  given  off),  (b)  of  nickelic 
hydroxide  from  nickelous  chloride  and  sodium  hypochlorite. 

5.  Tabulate  in  detail  the  chemical  relations  of  the  elements  cobalt 
and  nickel,  with  especial  reference  to  showing  the  resemblances  and 
differences. 


CHAPTER  XLIV 
THE    PLATINUM    METALS 

THE  remaining  elements  of  Mendelejeff's  eighth  group  divide*them- 
selves  into  two  sets  of  three  each.  Just  as  iron,  cobalt,  and  nickel 
have  similar  atomic  weights  and  much  the  same  specific  gravity 
(7.8-8.8),  so  ruthenium  (Ru,  at.  wt.  101.7),  rhodium  (Rh,  at.  wt.  103), 
and  palladium  (Pd,  at.  wt.  106.5)  have  specific  gravities  from  12.26 
to  11.5.  Similarly  osmium  (Os,  at.  wt.  191),  iridium  (Ir,  at.  wtf  193), 
and  platinum  (Pt,  at.  wt.  194.8)  form  a  triad  with  specific  gravities 
from  22.5  to  21.5.  Chemically,  ruthenium  shows  the  closest  resem- 
blance to  osmium,  and  both  are  allied  to  iron.  Similarly,  rhodium 
and  iridium,  and  palladium  and  platinum  are  natural  pairs. 

The  six  elements  are  found  alloyed  in  nuggets  and  particles  which 
are  separated  from  alluvial  sand  by  washing.  Platinum  forms 
60-84  per  cent  of  the  whole.  The  chief  deposits  are  in  the  Ural 
Mountains,  smaller  amounts  being  found  in  California,  Australia, 
Borneo,  and  elsewhere.  The  components  are  separated  by  a  com- 
plex series  of  chemical  operations. 

Ruthenium  and  Osmium.  —  These  metals  are  gray  like  iron, 
while  the  other  four  are  whiter  and  more  like  cobalt  and  nickel. 
They  also  resemble  iron  in  being  the  most  infusible  members  of  their 
respective  sets.  Both  melt  considerably  above  2000°.  They  like- 
wise resemble  iron  in  uniting  easily  with  free  oxygen,  while  the  other 
four* elements  do  not.  Ruthenium  gives  Ru02  and  even  Ru04, 
although  the  latter  oxide  is  more  easily  obtained  indirectly.  Osmium 
gives  OsO4,  "  osmic  acid,"  a  white  crystalline  body  melting  at  40° 
and  boiling  at  about  100°.  The  odor  and  irritating  effects  of  the 
vapor  re-c^l  chlorine  (Gk.,  ooywj,  odor).  The  substance  is  not  an 
acid,  or  even  an  acid  anhydride.  The  aqueous  solution  is  used  in 
histology,  and  stains  tissues  in  consequence  of  its  reduction  by 
organic  bodies  to  metallic  osmium.  It  is  affected  particularly  by 

508 


PLATINUM   METALS  509 

fat.  Osmic  acid  also  hardens  the  material  without  distorting  it. 
It  will  be  observed  that  ruthenium  and  osmium  have  a  maximum 
valence  of  eight. 

Rhodium  and  Iridium.  —  These  metals  are  not  attacked  by 
aqua  regia,  while  the  other  four  are  dissolved,  more  or  less  slowly. 
They  are  harder  than  platinum,  and  iridium  is  alloyed  with  this 
metal  for  the  purpose  of  increasing  its  resistance  to  the  action  of 
acids.  They  resemble  cobalt  in  having  no  acid-forming  properties. 
The  most  familiar  compounds  of  iridium  are  the  complex  chlorides 
X3IrCl0  (=  3XCl,IrCl3)  and  X2IrCl6  (=  2XCl,IrCl4).  The  solutions 
of  the  latter  are  red,  and  the  acid,  chloroiridic  acid  H2IrCle,  is  often 
found  in  commercial  chloroplatinic  acid  H2PtCle,  and  confers  upon 
it  a  deeper  color. 

Palladium  and  Platinum. —  Palladium  is  the  only  metal  of  this 
family  which  is  attacked  by  nitric  acid.  Palladium  and  platinum 
form  -ous  and  -ic  compounds  of  the  forms  PdX2  and  PdX4  respec- 
tively. The  oxides  PdO  and  PtO  and  corresponding  hydroxides  are 
basic.  When  quadrivalent,  the  metals  appear  chiefly  in  complex 
compounds,  like  H2.PtCl6,  H2.PdCl6,  in  which  the  metal  is  in  the 
anion.  Platinum  gives  also  platinates  derived  from  the  oxide  PtO2. 

Palladium.  —  This  metal,  named  from  the  planetoid  Pallas,  is 
noted  chiefly  for  its  great  tendency  to  absorb  hydrogen.  At  1000° 
it  takes  up  about  650  times  its  own  volume.  The  amount  absorbed 
varies  continuously  with  the  concentration  (pressure)  of  the  hydro- 
gen, although  not  according  to  a  uniform  rule,  and  the  case  is  there- 
fore regarded  as  being  one  of  solid  solution.  When  a  strip  of 
palladium  is  made  the  cathode  of  an  electrolytic  cell,  over  900 
volumes  of  hydrogen  may  be  occluded.  This  absorbed  hydrogen,  in 
consequence  of  the  catalytic  influence  of  the  metal,  reacts  more 
rapidly  than  does  the  gas,  and  consequently  a  strip  of  hydrogenized 
palladium  will  quickly  precipitate  copper  and  other  metals  less 
electro-positive  than  hyd^en  and  will  reduce  ferric  and  other 
reducible  salts: 

OuS04  +  H2  ->  H2S04  4-  Cu,    or    Cu"  4-  H2  ->  2H'  +  Cu. 
2FeCl3  +  H2  ->  2FeCl2  +  2HC1,    or    2Fe"*  +  H2  ->  2Fe"  +  2H. 


510  COLLEGE   CHEMISTRY 

Platinum.  —  This  metal  (dim.  of  Sp.  plata,  silver)  is  grayish- 
white  in  color,  and  is  very  ductile.  At  a  red  heat  it  can  be  welded. 
It  does  not  melt  in  the  Bunsen  flame  but  fuses  easily  in  the  oxyhy- 
drogen  jet.  On  account  of  its  very  small  chemical  activity  it  is  used 
in  electrical  apparatus  and  for  making  wire,  foil,  and  crucibles  and 
other  vessels  for  use  in  laboratories.  It  interacts  with  fused  alkalies, 
giving  platinates.  The  oxygen  acids  are  without  action  upon  it, 
but  the  free  chlorine  in  aqua  regia  converts  it  into  chloroplatinic  acid 
H2PtCV 

The  metal  condenses  oxygen  upon  its  surface  and  it  dissolves 
hydrogen.  The  finely  divided  forms  of  the  metal,  such  as  platinum 
sponge  made  by  igniting  ammonium  chloroplatinate  (NH4)2PtCl6, 
and  platinum  black  made  by  adding  zinc  to  chloroplatinic  acid,  show 
this  behavior  very  conspicuously.  They  cause  instant  explosion  of 
a  mixture  of  oxygen  and  hydrogen,  in  consequence  of  the  heat 
developed  by  the  rapid  union  of  that  part  of  the  gases  which  is  con- 
densed in  the  metal.  A  heated  spiral  of  fine  platinum  wire  will 
continue  to  glow  if  immersed  in  the  mixture  of  alcohol  vapor  and 
oxygen  formed  by  leading  oxygen  through  liquid  alcohol.  The  heat 
is  developed  by  the  interaction  which  takes  place  between  the  sub- 
stances with  great  speed  at  the  surface  of  the  platinum.  Platinum 
sponge  is  the  active  constituent  of  the  contact-mass  used  in  making 
sulphur  trioxide  (p.  257). 

Platinum  is  the  only  otherwise  suitable  substance  which  has  the 
same  coefficient  of  expansion  as  glass,  and  it  is  consequently  fused 
into  incandescent  bulbs  and  furnishes  the  electrical  connection  with 
the  filament  in  the  interior.  Large  amounts  are  also  consumed  in 
photography.  The  price  of  the  metal  is  subject  to  great  variations, 
since  a  rainy  season  in  the  Caucasus  will  render  larger  amounts 
accessible  to  the  miners;  but,  on  the  whole,  the  many  applications 
which  have  been  found  for  it  have  tripled  its  price  in  the  last  twenty 
years. 

When  special  resistance  to  chemical  or  mechanical  influences  is 
required,  as  in  standard  meters  for  international  reference,  or  points 
of  fountain  pens,  the  alloy  with  iridium  is  employed. 

Compounds  of  Platinum.  —  Platinous  chloride  is  made  by  pas  > 
ing  chlorine  over  finely  divided  platinum  at  240-250°  or  by  heating 
chloroplatinic  acid  to  the  same  temperature.  It  is  greenish  and 


PLATINUM  METALS  511 

; 

insoluble  in  water,  but  forms  with  hydrochloric  acid  the  soluble 
chloroplatinous  acid  H2PtCl4.  Potassium  chloroplatinite  K2PtCl4  is 
used  in  making  platinum  prints.  Bases  precipitate  black  platinous 
hydroxide  Pt(OH)2  which  interacts  with  acids  but  not  with  bases. 
Gentle  heating  gives  the  oxide  PtO  and  stronger  heating  the  metal. 
With  potassium  cyanide  and  barium  cyanide  soluble  platino-cyanides, 
K2Pt(CN)4,3H2O  and  BaPt(CN)4,4H20,  are  formed.  These  sub- 
stances, when  solid,  show  strong  fluorescence,  converting  X-rays  as 
well  as  ultra-violet  rays  into  visible  radiations.  The  barium  salt  is 
used  to  coat  screens  on  which  the  shadows  cast  by  X-rays  are 
received. 

Chloroplatinic  acid  H2PtCl6,6H20  is  made  by  treating  the  metal 
with  aqua  regia,  and  forms  reddish-brown  deliquescent  crystals. 
With  potassium  and  ammonium  salts,  it  yields  the  sparingly  soluble, 
yellow  chloroplatinates  K2PtClQ  and  (NH4)2PtCl6  (cf.  p.  369),  in 
solutions  of  which  the  platinum  migrates  towards  the  anode  and 
silver  salts  precipitate  Ag2PtCle  and  not  silver  chloride.  'Bases 
interact  with  chloroplatinic  acid,  giving  a  yellow  or  brown  precipitate 
of  platinic  hydroxide  Pt(OH)4.  This  substance  interacts  with  bases 
to  give  platinates,  like  Na2H10Pt3O12,H2O.  Both  sets  of  platinum 
compounds  interact  with  hydrogen  sulphide,  giving  the  sulphides, 
PtS  and  PtS2  respectively.  These  are  black  powders  which  dissolve 
in  yellow  ammonium  sulphide  solution,  much  as  do  the  sulphides  of 
gold,  arsenic,  and  other  metals,  giving  ammonium  sulphoplatinates. 


INDEX. 


***  Acids  are  all  listed  under  "  acid  "-  and  salts  under  the  positive  radical. 


ACETONE,  318 
Acetylene,  136,  330,  341 
Acid,  acetic,  64,  318,  331,  382 

antimonic,  467,  469 

arsenic,  463,  465 

arsenious,  465 

boracic,  349 

boric,  349 

bromic,  197 

carbolic,  294 

carbonic,  323 

chloric,  195 

chloroiridic,  509 

chloroplatinic,  299,  509,  510,  511 

chlorous,  196 

chromic,  479 

disulphuric,  263 

formic,  325,  330 

hydrazoic,  280,  28-* 

hydriodic,  167,  168,  see  Hydrogen 
iodide 

hydrobromic,    163,    see    Hydrogen 
bromide 

hydrochloric,  63,  122,  see  Hydrogen 

chloride 
fractions  ionized,  227 

hydrocyanic,  335 

hydrofluoboric,  349 

hydrofluoric,  170. 

hydrofluosilicic,  345 

hypobromous,  197 

hypochlorous,  115,  190 
bleaching  by,  192,  194 

hyponitrous,  300 

hypophosphoric,  309 

hypophosphorous,  309 

hyposulphurous,  258 

iodic,  198 

lactic,  333 

metaphosphoric,  309,  312 

metastannic,  454 

muriatic,    123,    see    Acid,    hydro- 
chloric 

nitric,  292,  301 
fuming,  294,  296 


Acid,  nitric,  oxidizing  actions  of,  297 

test  for,  296 
nitrosylsulphuric,  259 
nitrous,  299 

orthophosphoric,  309,  311 
osmic,  508 
oxalic,  325,  331 
palmitic,  334 
perchloric,  196,  197 
periodic,  198 
permanganic,  491 
persulphuric,  267 
phosphoric,  118,  161 
phosphorous,  309,  313 
picric,  294 
prussic,  335 

pyrophosphoric,  309,  312 
salts,  231 
selenic,  270 
silicic,  346 
a-stannic,  454 
/?-stannic,  454 
sulpharsenic,  466 
sulphuric,  63,  257,  259,  264 

chamber  process,  259 

contact  process,  257 

electrolysis  of,  216 

heat  of  solution,  100 

properties,  263,  264 

uses,  264 

sulphurous,  257,  265 
sulphydric,  252 
tannic,  449,  499 
telluric,  270 
thiosulphuric,  266 
Acidimetry,  240 
Acids,  active,  229 
activity  of,  229 
chemical  behavior  of,  201 
common,  63 

dibasic,  ionization  of,  252 
feeble,  229 

interaction  with  oxides,  111 
ionic  substances  from,  232 
monobasic,  dibasic,  etc.,  231 


513 


514 


INDEX 


Acids,  nomenclature  of,  186 

non-ionic  formation  of,  246 

of  constant  boiling-point,  120 

organic,  330 

phosphoric,  308 
tests  for,  313 

polythionic,  258 

weak,  229 

xanthoproteic,  294 
Actions,  balanced,  176 

non-ionic,  246 

reversible,  174 

completion  of,  184 
Activity,  apparent,  177 

chemical,  18 

intrinsic,  177 

of  ionogcns,  229 

of  metals,  order  of,  244,  245     . 

see  Chemical  activity 
Adsorption,  80,  319,  449 
Affinity,  apparent  and  intrinsic,  177 

chemical,  120 

constant,  181 

see  Chemical  activity 
Agate,  346 
Agar-agar,  218 
Air,  286 

composition,  288 

fixed,  322 

liquid,  290 
Albite,  373 
Albumins,  279,  280 
Alcohol,  absolute,  333 

ethyl,  329,  333,  478 

methyl,  318,  333 
Aldehyde,  478 
Alizarine,  449 
Alkalies,  66,  81,  232 

metals  of,  360 
Alkalimetry,  240 
"Alkaline  air, "280 
Alkaline  earths,  metals  of,  388 
Allotropic  forms,  306 
Alloys,  352,  411 
Alum,  burnt,  447 

chrome,  501 

common,  447 

ferric,  501 

neutral,  447 
Alumina,  446 
Aluminates,  445 
Aluminium,  443 

bronze,  444 

carbide,  320..  329,  358 

chloride,  445 

hydroxide,  445 

***  Acids  are  all  listed  under  "  acid  ' 


Aluminium,  -ion,  450 

oxide,  446 

sulphate,  446 

sulphide,  447 
Aluminothermy,  355,  444 
Alums,  447 
Amalgam,  ammonium,  371 

sodium,  373 
Amalgams,  352,  436 
Amblygonite,  380 
Ammonia,  281 
Ammonia-soda  process,  378 
Ammonio-argentic-ion,  419 

-cupric  compounds,  412,  413,  414 
416 

-cuprous  compounds,  412,  413,  41- 
Ammonium  amalgam,  371 

bicarbonate,  370 

bisulphate,  370 

bi  tart  rate,  371 

carbonate,  370 

chloride,  283,  370 
heat  of  solution,  100 

chloroplatinate,  371,  511 

compounds,  test  for,  283 

dichromate,  481 

-hydrogen  sulphide,  370 

hydroxide,  282,  283 

-ion,  371 

molybdate,  482 

nitrate,  300,  301,  370 

nitrite,  279 

persulphate,  370 

polysulphides,  371 

sulphate,  370 

sulphide,  370 
yellow,  371 

sulphantimoniate,  469 

sulpharsenate,  466 

sulphostannate,  455 

thiocyanate,  336,  370 
Amorphous  bodies,  249 
Amperemeter,  22G 
Analysis,  gravimetric,  242 

qualitative  (cations),  439 

volumetric,  242 
Analytical  reactions,  aluminium,  450 

ammonium,  371 

barium,  406 

cadmium,  435 

calcium,  403,  406 

chromium,  482 

cobalt,  506 

copper,  416 

iron,  502 

lead,  460 
and  salts  under  the  positive  radical. 


INDKX 


515 


Analytical  reactions,  magnesium,  lo 

manganese,  492 

mercury,  439 

nickel,  506 

potassium,  369 

silver,  422 

sodium,  380 

strontium,  404,  406 

tin,  455 

zinc,  434 
Anglesite,  459 
Anhydride,  chromic,  477,  479 

hypochlorous,  191 

iodic,  198 

nitric,  294 

nitrous,  260,  300 

perchloric,  196 

permanganic,  491 

phosphoric,  51,  308 

phosphorous,  308 

sulphuric,  257 

sulphurous,  51,  256 
Anhydrides,  51 

relation  to  acids,  310 
Anhydrite,  399 
Animal  charcoal,  319 
Anion,  224 
Anode,  224 
Anthracite,  320 
Antichlor,  380 
Antimony,  467 

Chlorides,  468 

halides,  468 

nitrate,  468 

oxides,  467 

oxychloride,  468 

sulphate,  467,  468 

sulphides,  469 

trichloride,  468 

tetroxide,  467 

trioxide,  467 

trisulphide,  467 
Antimonyl,  468 
Apatite,  303,  401 
Aq,  meaning  of  symbol,  52 
Aqua  regia,  299,  302 
Aqueous  tension,  60 
Argentic  compounds,  see  Silver 

-ion,  422 
Argon,  286,  290 

Arithmetical  problems,  43,  60,  143 
Arrows,  vertical,  118 
Arsenic,  462 

family,  471 

pentasulphide,  466 

pentoxide,  465 


Arsenic,  sulphides,  466 

trichloride,  464 

trioxide,  464 

white,  464 
"  Arsenic,  "464 
Arsenious,  see  Arsenic  tri- 
Arsenuretted  hydrogen,  463 
Arsine,  463 
Asbestos,  428 
Ash,  black,  376 

bone,  401 
Assaying,  425 
Atakamite,  411 
Atmosphere,  286 
Atomic  heats,  135 
Atomic  hypothesis,  89,  151 
Atomic  weight,  36,  130 

of  a  new  element,  147 
Atomic    weights,    advantages    over 
equivalents,  134 

determination  of,  277 

table  of,  inside  rear  cover 
Atoms,  127 

properties  of,  152,  156 
Attributes  of  bodies,  23 
Aurates,  425 
Avogadro's  hypothesis,  89,  125 

BACTERIA,  331,  333 

nitrifying,  280 
Balanced  actions,  176 
Barff's  process,  498 
Barite,  248,  405 
Barium,  405 

carbide,  358 

carbonate,  405 

chlorate,  406 

chloride,  405 

chromate,  479 

compounds  of,  405 

dioxide,  see  Peroxide 

hydroxide,  406 

-ion,  406 

nitrate,  406 

oxide,  45,  406 

peroxide,  45,  211 

sulphate,  405 

sulphide,  405 
Barometer,  58 

correction  of,  59 
Baryta-water,  406 
Bases,  81 

activity  of,  229 

chemical  behavior  of,  201 

ionic  substances  from,  232 

non-ionic  formation  of,  246 


***  Acids  are  all  listed  under  "  acid  "  and  salts  under  the  positive  radical. 


516 


INDEX 


Bases,   solubilities   of,   Table    inside 

front  cover 
Basic  salts,  231 
Basicity,  231 
Bauxite,  21,  45 
Bead  tests,  312,  350 
Bell-metal,  411 
Bengal  lights,  466 
Benzine,  328 
Beryl,  348,  428 
Beryllium,  427 
Bessemer  process,  497 
Bicarbonates,    see   Hydrogen   carbo- 
nates 

Bichromates,  see  Bichromates 
Binary  compounds,  70 
Bischofite,  108 
Bismuth,  470 

compounds,  470,  471 
Bisulphates,  265,  see  Hydrogen  sul- 
phates 

Bisulphides,  253 

Bisulphites,  see  Hydrogen  sulphites 
Black-ash,  376 
Black-lead,  318 
Blast  furnace,  495 
Bleaching,  192,  193,  210,  213,  257 

powder,  189,  399 
Blood  charcoal,  319 
Blue-stone  (blue  vitriol),  415 
Body,  definition  of,  22 
Boiling-points  of  solutions,  abnormal, 

207 

Bone-ash,  303,  401 
Bone  black,  319 
Bones,  distillation  of,  331 
Boracite,  350 
Borates,  350 
Borax,  348,  350,  380 
Boron,  348 

compounds  of,  349 

nitride,  280 
Boyle's  law,  59,  87 
Brass,  411 
Braunite,  487 

Breathing,  chemistry  of,  52 
Bricks,  449 
Brimstone,  248 
Brin's  oxygen  process,  46,  391 
Britannia  metal,  452 
Bromine,  158 
Bronze,  411,  444 
Brucite,  429 
Bunsen  flame,  339,  341 
Butane,  327 

*#*  Acids  are  all  listed  under  "  acid 


CADMIUM,  434 

compounds,  434 

-ion,  435 
Caesium,  369 
Calamine,  431 
Calcination,  256 
Calcite,  94,  391 
Calcium,  389 

aluminate,  446 

bicarbonate,  391 

bisulphite,  266,  332 

carbide,  321,  390 

carbonate,  94,  322,  390 

chlorate,  195 

chloride,  390 

chromate,  479 

fluoride,  168,  390 

hydride,  390 

-hydrogen  sulphite,  266,  332 

hydroxide,  393 

-ion,  403 

light,  338 

manganite,  490 

nitride,  280,  281,  390 

oxalate,  331,  394 

oxide,  322,  392 

phosphate,  303 

phosphates,  401 

phosphide,  307 

silicate,  402 

sulphate,  95,  399 

sulphide,  376,  400 

hydrolysis  of,  254 
Calculations,  43 
Calomel,  436,  439 
Calorie,  55 
Camphor,  268 
Cane-sugar,  332 

Carbides,  characteristics  of,  358 
Carbohydrates,  332 
Carbon,  316 

amorphous,  318 

chemical  relations  of,  316 

classes  of  compounds,  327 

hydrogen  compounds,  327 

properties,  319,  320 

dioxide,  321 

disulphide,  326 

monoxide,  325 

compounds  with  metals,  502,  506 
poisoning,  326 

pure,  319 

tetrachloride,  329 
Carbonates,  323,  324 

acid,  see  Hydrogen  carbonates 

characteristics  of,  358 

"  and  salts  under  the  positive  radical. 


INDEX 


517 


Carbonates,  effect  of  heating,  322 

Carbonyl  chloride,  320 

Carborundum,  320,  344 

Carnallite,  362,  428 

Cassiterite,  452 

Casts,  400 

Catalytic  action,  54 

Catalytic  actions,  74,  114,  161,  257, 

268,  299,  302,  332,  336 
Cathode,  224 
Cations,  224 
Cat's  eye,  346 
Cause,  definition  of,  19 
Caustic  potash,  232 
Caustic  soda,  232,  374 
Celestite,  248,  404 
Cellulose,  318,  325,  332 
Cement,  393 
Cerite,  460 
Cerium,  443,  460 

dioxide,  338 
Chalcocite,  409 
Chalk,  391 

Chamber  process,  259 
Chance's  process,  377 
Characteristics  of  chemical  phenom- 
ena, 3,  11,  12,  18,  29,  33 
Charcoal,  318,  319,  320  • 
Charles' law,  60,  88 
Chemical  action,  means  of  initiating, 
53 

actions,  reversible,  174 

activity,  18 
cause  of,  19 
measured   by  fractions  ionized, 

229 
measurement  of,  19 

change,   extent  of  with  ionogens, 

varieties  of,  9,  68,  81,  115,  124, 
243  (ionic) 

changes,  simultaneous,  194,  196 

energy,  17 

equilibrium,  see  Equilibrium 

phenomena,  illustrations  of,  6,  7,  8 

properties,  specific,  75 

relations  of  elements,  116,  157 
Chemistry,  methods  of  work  in,  24 
Chinese  white,  432 
Chitin,  391 
Chlorates,  195 
Chlorides,  characteristics  of,  357 

electrolysis  of,  109 

interaction  with  acids,  117 

modes  of  preparing,  123 

solubility,  122 


Chlorine,  108 

dioxide,  196 

heptoxide,  197 

monoxide,  191 

test  for,  114 

uses,  116 
Chloroform,  329 
Chlorophyll,  324,  495 
Chrome-alum,  481 
Chrome- red,  479 
Chromic,  anhydride,  477,  479 

chloride,  480 

hydroxide,  480 

-ion,  482 

oxide,  481 

sulphate,  481 
Chromite,  475 
Chromites,  480 
Chromium,  474 

carbide,  358 

trioxide,  481 

Chromous  compounds,  481,  482 
Cinnabar,  435,  438 
Clay,  348,  449 
Coal,  320 
Coal  gas,  340 
Cobalt,  503 

aluminate,  450 

complex  compounds,  504 

zincate,  433 

Cobaltic  compounds,  504 
Cobaltite,  503 
Cobaltous  compounds,  503 
Coke,  318,  320 
Colemanite,  348 
Colloidal  solution,  96,  500 
Columbium,  472 
Combination,  9 

Combining  proportions,  measure- 
ment of,  29 
Combining  weights,  31 

law  of,  34 
Combustion,  51 

heat  of,  55 

spontaneous,  56 

Common  ion,  effect  of  adding,  386 
Complex  ions,  formation  of,  413 
Components,  physical,  22 
Compound,  7 
Compounds,  binary,  70 

unsaturated,  257 
Concentration,  molecular,  178 

molecular,  law  of,  180 

of  gases,  88,  179 
Conditions,  24 


***  Acids  are  all  listed  under  "  acid  "  and  salts  under  the  positive  radical. 


518 


INDEX 


Conductivity,  electrical,  of  metals,  353 

electrical,  of  solutions,  220 

interpretation  of,  226 
Condy's  disinfecting  fluid,  491,  492 
Congo  red,  243 
Copse rvation  of  energy,  17 

of  mass,  12 
Constant,  equilibrium,  181 

ion-product,  385 

ionization,  225 

molecular  depression,  204 
Constituents,  chemical,  22 
Contact,  actions,  see  Catalytic 

process,  257 
Copper,  408  *  % 

pyrites,  248 

see  Cupric  and  Cuprous 
Copperas,  499 
Coral,  391 

Corrosive  sublimate,  436 
Cotton,  332,  448 
Crayon,  391 
Crocoisite,  475 
Cryolite,  443 
Crystal,  forms,  93 
Crystallization,  water  of,  82,  84,  see 

Hydrates 

Cullinan  diamond,  317 
Cupel,  417,  425 
Cupellation,  418 
Cupric  acetate,  415 

ammonio-,   compounds,   412,    41:*, 
414,  416 

bromide,  235,  381,  412 

carbonate  (basic),  414 

chloride,  411 

chloride  (basic),  411 

cyanide,  335 

cyanides,  414 

ferrocyanide,  416,  502 

hydroxide,  414 

iodide,  412 

-ion,  416 

nitrate,  414 

electrolysis  of,  216 

oxide,  414 

oxychloride,  411 

phosphide,  307 

sulphate,  415 

electrolysis  of,  218 
hydrated,  82/83,  415 

sulphide,  416 
Cuprite,  409 

Cuprous,  ammonio-,  compounds,  412, 
413,  414 

bromide,  412 

*#*  Acids  are  all  listed  under  "acid 


Cuprous,  chloride,  412,  413 

cyanide,  335,  414 

iodide,  413 

-ion,  416 

oxide,  413 

sulphide,  416 
Cyanogen,  335 

D  ALTON'S  law  of  partial  pressures,  60 

Deacon's  process,  110 

Decane,  328 

Decantation,  26 

Decomposition,  10 

Decrepitation,  366 

Definite  proportions,  law  of,  29 

Degrees  of  ionization  (data),  228 

Deliquescence,  364 

Densities  of  gases,  126 

measurement  of,  61 
Depression  constant,  molecular,  204 
Depression,  freezing-point,  204 

abnormal,  205 
Dextrose,  332,  333 
Dialysis,  500 
Diamond,  317 
Diaspore,  445 
Diatomic  molecules,  138 
Dichromates,  476 

properties  of,  477 
Diffusion,  in  gases,  74 

in  solution,  101 
Dimorphous  substances,  249 
Disinfectants,  257,  399,  491 
Displacement,  68,  75 

ionic,  243 
Displacement  of  equilibria,  183 

ionic,  234 
Dissociation,  81 

hydrolytic,  see  Hydrolysis 

in  solution,  201,  207 

measured  by  density,  146 

pressure,  391 
Distillation,  26 

fractional,  328 

under  reduced  pressure,  197,  211 
Disulphates,  263 
Dog-tooth  spar,  391 
Dolomite,  317 
Double  decomposition,  10 

in  solution,  236 
Double  salts,  231 
Dulong  and  Petit's  law,  135 
Dust,  in  air,  286,  287 
Dutch  metal,  114 
Dyeing,  448 
and  salts  under  the  positive  radical. 


INDEX 


519 


EARTHENWARE,  449 
Earths,  alkaline,  388 

rare,  442 
Efflorescence,  83 
Egg-shells,  391 

Electric  furnace,  304,  320,  321 
Electrical    conductivity,    of   electro- 
lytes, 226 

of  metals,  353 

Electricity  in  chemical  change,  13 
Electrogilding,  425 
Electrolysis,  215 

explanation  of,  220 

of  chlorides,  109 

of  dilute  acids,  64 

primary  products  of,  216 

secondary  products  of,  216 
Electrolytes,  214 

Electrolytic  deposition,  copper,  410 
Electromotive,  series,  243,  245 

applications  of,  66,  76 
Electrons,  485 
Electroplating,  copper,  410 

gold,  425 

nickel,  505 

silver,  421 
Element,  definition  of,  20 

decomposition  of,  21,  485 
Elementary     substances,     liberation 
of,  67 

molecular  weights  of ,  137 
Elements,  base-forming,  see  Elements 
metallic 

chemical  relations  of,  116,  157 

metallic,  233,  271,  353  ' 

metallic  and  non-metallic,  82,  276 

non-metallic,  233,  272 

positive  and  negative,  224 
,     quantities  of  interrestrial  material, 
21 

similar,  definition,  157 
Emeralds,  428 
Endothermal  actions,  55 
Energy,  15 

conservation  of,  17 

internal,  17 
Enzymes,  333 
Epsom  salts,  430 
Equations,  40 

making  of,  41,  50,  159,  297 

molecular,  applications  of,  143 
interpretation  of,  142 
making  of,  142 

partial,  159 
Equilibrium,  characteristics  of,  91 

chemical,  174 


Equilibrium,  chemical,  characteristics 

of,  176 

displacement  of,  183 
constant,  181 
in  saturated  solution,  105 
ionic,  224,  381 

considered  quantitatively,  381 
displacement  of,  234 
in  saturated  solutions,  384 
physical,  91 

Equivalent  weights,  32,  134 
Esters,  334 
Ethane,  318,  327,  328 
Ethyl  acetate,  334/ 
Ethylene,  329,  3*1 

bromide,  330* 
Evaporation,  '26 
Exothermal  actions,  55 
Explanation,  its  nature,  5 
Explosives,  301 
Extraction,  103 

FATS,  334 
Feldspar,  348,  443 

soda,  373 

Fermentation,  331,  332 
Ferrates,  494 
Ferric  alum,  501 

chloride,  500 

compounds,  .500 

hydroxide,  500  ' 

-ion,  502 

oxide,  494,  500 

sulphate,  501 

sulphide,  501 

thiocyanate,  336,  502 
Ferrous  carbonate,  499 

chloride,  498 

hydroxide,  499 

iodide,  362 

-ion,  502 

oxide,  499 

sulphate,  296,  499 

sulphide,  7,  499 
Fertilizers,  368,  401,  497 
Filters,  charcoal,  319 

Pasteur,  78 
Fireproofing,  380,  483 
"  Fixed  air,"  322 
Flame,  337 

Flash-light  powder,  428 
Flint,  346 
Fluorine,  168 
Fluorite,  168,  390 
Fluor-spar,  390 
Flux,  355 


***  Acids  are  all  listed  under  "  acid  "  and  salts  under  the  positive  radios!. 


520 


INDEX 


Fools'  gold,  501 
Formaldehyde,  257 
Formulae,  39 

graphic,  301 

making  of,  40,  50 

molecular,  of  compounds,  133,  136 
of  simple  substances,  137,  139 

reaction,  83 

structural,  199 

see  Equations 

Formulation,  excess  of  one  ion,  382 
Fractions  ionized  (data),  228 
Franklinite,  431,  500 
Freezing  mixtures,  204,  282,  390 
Freezing-points,  abnormal,  205 

of  solutions,  204 
Fulminating  mercury,  438 
Furnace,  blast,  495 

electric,  304,  320,  321 

reverberatory,  375,  376 
Fusel  oil,  319 
Fuses,  detonating,  195 

G.  M.  V.,  128 

Gahnite,  446 
Galena,  see  Galenite 
Galenite,  248,  417,  456 
Gall,  334 
Gallium,  277,  442 
Garnet,  94 
Garnierite,  505 
Gas,  coal,  340 

illuminating,  340 

oil,  340 

water,  325,  340 
Gases,  densities  of,  61,  126 

drying  of,  390 

liquefaction  of,  289 

mixed,  60 

relative  volumes  of,  145 

relative  weights  of,  61,  145 

solubility  of,  103 
Gasolene,  328 
Gay-Lussac's  law,  see  Charles'  law 

of  combining  volumes,  84,  89 
Generalization,  definition  of,  4 
Germanium,  277,  451 
German  silver,  411,  505 
Glass,  402 

cobalt,  504 

colored,  403 

etching  of,  171 
Glauber's  salt,  379 

solubility  of,  106 
Glucinum,  427 
Glucose,  332,  414 


Glycerine,  294,  334 
Gold,  422 

compounds  of,  424 

extraction  of,  116 

fools',  501 

Goldschmidt's  process,  444 
Gram-molecular  volume,  128 
Granite,  348 
Grape-sugar,  332 
Graphite,  318 

artificial,  109 
Gravimetric,  242 
Gun-cotton,  294,  301 
Gun-metal,  411 
Gunpowder,  26,  366 
Gypsum,  95,  248,  399 

HEMATITE,  494 

Haemoglobin,  52,  326,  495 

Halite,  108 

Halogen  family,  171 

Halogens,  chemical  relations  of,  158, 

172,  199 

Hardness  (water),  77,  392 
Heat,  of  combustion,  55 

of  fusion,  78 

of  reaction,  13,  55 

of  solution,  100 

of  vaporization,  80 
Heating,  256 
Heavy-spar,  405 
Helium,  290,  485 
Henry's  law,  103 
Heptane,  327,  328 
Hexacontane,  327 
Hexadecane,  327,  328 
Hexane,  327,  328 
Homogeneous  systems,  181 
Horn-silver,  417 
Hydrargyllite,  445 

Hydrates      (substances      containing 
water  of  crystallization),  82, 84 
Hydrazine,  280,  283 
Hydrocarbons,  327 

saturated,  327 

unsaturated,  329 
Hydrogen,  63 

active  (nascent),  302 

arsenu  retted,  463 

chemical  properties,  74 

nascent,  302 

occurrence,  63 

physical  properties,  72 

preparation,  64,  65,  66,  67 
Hydrogen  bromide,  161 

see  Acid,  hydrobromic 


***  Acids  are  all  listed  under  "  acid  "  and  salts  under  the  positive  radical. 


INDEX 


521 


Hydrogen  chloride,  117,  387 

composition,  121 

preparation,  theory  of,  118 

see  Acid,  hydrochloric 
Hydrogen  fluoride,  170 
Hydrogen  iodide,  166 

dissociation  of,  174,  180,  183 
Hydrogen-ion,  232 
Hydrogen,  peroxide,  210 

selenide,  269 

sulphate,  263 

sulphide,  250 

see  Acid,  sulphuric 
Hydrolite,  390 
Hydrolysis,  163 

of  salts,  353 
Hydrosulphate-ion,  264 
Hydrosulphide-ion,  252 
Hydroxide-ion,  232 
Hydroxylamine,  280 
Hypo,  379 
Hypochlorites,  188 
Hypochlorous  anhydride,  191 
Hypothesis,  atomic,  89,  151 

Avogadro's,  89,  125 

kinetic-molecular,  86 

of  ions,  219 

ICE,  78 

Iceland  spar,  391 
Identification,  means  of,  24 
Illuminating-gas,  340 
Illuminating  lamps,  Nernst,  460 

Welsbach,  338 
Impure,  definition  of,  22 
Indicators,  242 
Indigo,  192,  194,  210 
Indium,  442 
Ink,  marking,  420 

printers',  318 

sympathetic,  503 

writing,  499 

Insoluble  salts,  theory  of  precipita- 
tion, 394,  398 

theory  of  solution,  395,  396 

solution  to  give  complex  ions,  413 
Internal  energy,  17 
Internal  rearrangement,  10 
lodic  anhydride,  198 
Iodine,  164 

atomic  weight  of,  277 

chlorides,  172 

pentoxide,  198 

tincture  of,  166 

union  with  hydrogen,  166,  174,  180, 
183 

***  Acids  are  all  listed  under  "acid 


Iodine,  vapor,  density  of,  147,  166 

lodoform,  329 

Ion,  common,  effect  of  adding,  386 

-product  constant,  385 
Ionic  equilibrium,  224 

considered  quantitatively,  381 

displacement  of,  234 

in  saturated  solutions,  384 

with  single  ionogen,  233 
Ionic  substances,  231 
;    names  of,  224 
'Ionium,  485 
lonization,  214 

constant,  225 

degrees  of  (data),  228 

hypothesis  of,  219 

repression  of,  386 
lonogens,  223 

classes  of,  231 

non-ionic  formation  of,  245 
Ions,  charges  on,  220 

difficulties  presented  by  hypothesis 
of,  220 

hypothesis  of,  219 

migration  of,  217 

nomenclature  of,  223 

speed  of  migration,  219 
Iridium,  508,  509 
Iron,  494 

alum,  501 

carbide,  496 

carbonyls,  502 

cast,  496 

chemical  properties,  498 

cyanides,  501 

dialysed,  500 

galvanised,  431 

magnetic  oxide,  494,  501 

meteoric,  494 

ore,  bog,  500 

ore,  brown,  500 

pyrites,  494 

spiegel,  496 

-stone,  494 

wrought,  496 

see  Ferrous  and  Ferric 
Isatin,  192,  194 

JASPER,  346 
Jute,  332 

KAINITE,  368 
Kaolin,  348,  443,  449 
Kelp,  164 
Kerosene,  328 
Kinetic-molecular  hypothesis,  86 

and  salts  under  the  positive  radical. 


522 


INDEX 


Kohinoor  diamond,  317 
Krypton,  291 

LAMPBLACK,  318 
Lanthanum,  443 
Laughing  gas,  301 
Law,  Boyle's,  59,  87 

Charles',  60,  88 

definition  of,  4 

Dulong  and  Petit's,  135 

Gay-Lussac's,  84,  89 

Henry's,  103 

of  combining  weights,  34 

of  diffusion,  74 

of  freezing-point  depression,  204 

of  ion-product,  385 

of  molecular  concentration,  180 

of  octaves,  272 

of  partition,  103 

periodic,  276 
Lead,  456 

acetate,  459 

carbonate,  459 

chlorides,  457 

chromate,  460,  479 

dioxide,  111 

hydroxide,  458 

iodide,  457 

nitrate,  459 

oxides,  457 

pencils,  318 

red,  458 

sugar  of,  459 

sulphate,  459 

sulphide,  460 

white,  459 

Le  Blanc  soda  process,  375 
Lepidolite,  380 
Levulose,  333 
Light,  from  chemical  action,  305 

in  chemical  action,  9,  113,  191,  324 

see  Photochemistry 
Lignin,  318 
Lime,  building,  393 

burning,  392 

chloride  of,  see  Bleaching  powder 

in  water,  324 

-light,  338 

superphosphate  of,  402 

-water,  393 
Limestone,  390 
Linen,  332,  448 
Liquefaction  of  gases,  289 
Liquid  air,  290 
Litharge,  457 

***  Acids  are  all  listed  under  "  acid 


Lithium,  380 

carbide,  358 

compounds  of,  381 

nitride,  280 
Litmus,  242,  243 
Lodestone,  501 
Luminosity,  cause  of,  341 
Lunar  caustic,  420 
Lyes,  232 

MACARTHUR-FOREST  process,  423 
Magnalium,  444 
Magnesia  alba,  430 
Magnesia,  calcined,  429 
Magnesite,  428 
Magnesium,  428 

ammonium  phosphate,  311,  430 

carbonate,  428,  430 

chloride,  428 

hydroxide,  429 

-ion,  430 

nitride,  280,  281,  428 

oxide,  429 

sulphate,  430 

sulphide,  430 
Magnetite,  494 
Malachite,  409,  414 
Malleability,  order  of,  351 
Manganates,  487,  490 
Manganese,  487 

carbide,  358 

classes  of  compounds,  487 

dioxide,  111 
catalysis  by,  47 

oxides  of,  488,  491 
Manganic  compounds,  487,  490 
Manganites,  487,  490 
Manganous  compounds,  487,  489 
Marble,  391 
Marsh  gas,  328 
Marsh's  test,  463 
Mass,  action,  180 

conservation  of,  12 
Matches,  306 
Matrix,  248 

Mechanical  energy,  in  chemistry,  15 
Meerschaum,  428 
Mendelejeff's  table,  274 
Mercuric  chloride,  436 

fulminate,  438 

-ion,  439 

oxide,  7,  437 
Mercurous-ion,  439 
Mercury,  435 

-ammonia  compounds,  438 

and  salts  under  the  positive  radical. 


INDEX 


523 


Mercury,  chlorides,  436 

iodides,  437 

nitrates,  438 

oxides,  437 

sulphides,  438 

Metallic  elements,  82,  233,  271,  353 
Metals,   electrical   conductivities   of, 
353 

extraction  from  ores,  35.5 

found  native,  245,  354 

heavy  and  light,  351 

hydroxides  of,  356 

melting  points  of,  352 

native,  245,  354 

oxides  of,  356     , 

physical  properties  of,  351 

see  Metallic  elements 
"  Metals,"  recognition  of,  in  analysis, 

439 

Meteorites,  494 
Methane,  318,  327,  328 
Methyl,  orange,  243 
Mica,  348,  443 

lithia,  380 

Microcosmic  salt,  311,  312 
Migration,  ionic,  217 
Minium,  458 
Mirrors,  421 
Mixed  gases,  60 

salts,  231 
Mixture,  22 
Molar,  volume,  128 

solutions,  99 

weight,  128 
Molasses,  332 
Molds,  333 
Mole,  128 
Molecular,  concentration,  law  of,  180 

depression  constant,  204 

equations,  142 

applications  of,  143 

formulae,  133,  136 

hypothesis,  86 

weight,  205 

in  solution,  205,  212 

weights,  127 

of  elements,  137,  146,  147 
Molecule,  definition,  87 
Molecules,  diatomic,  138 

relative  weights  of,  125 
Molybdenum,  482 

carbide,  358 

compounds  of,  482,  483 
Mordants,  448,  449 
Mortar,  393 
Multiple  proportions,  law  of,  33 

***  Acids  are  all  listed  under  "  acid 


NAPHTHA,  328 

Nascent  hydrogen,  283,  302 

Negative  elements,  224 

radicals,  64 
Neodymium,  443 
Neon,  291 
Nernst  lamp,  460 
Neutral  salts,  231 
Neutralization,  189 

theory  of,  239 
Niccolite,  505 
Nickel,  505 

carbonyl,  506 

compounds,  505,  506 

glance,  505 

meteoric,  494 

plating,  505 

sulphate,  94 
Nitrates,  295 
Nitric  anhydride,  294 

oxide,  295 
Nitrides,  280 
Nitrites,  299 
Nitro-derivatives,  294 
Nitrogen,  279 

chloride,  284 

compared    with    phosphorus    and 
sulphur,  314 

dioxide,  see  Nitric  oxide 

family,  471 

iodide,  284 

monoxide,  300 

oxides,  292 

oxygen  acids  of,  292 

tetroxide,  296 

trioxide,  260,  300 
Nitroglycerine,  294,  301 
Nitrosyl,  259 

chloride,  299 
Nitrous  anhydride,  260,  300 

oxide,  300 
Nomenclature,  51,  186 

ionic  substances,  223 
Nonane,  328 
Non-electrolytes,  214 
Non-ionic  actions,  245 
Non-metallic  elements,  82,  233,  272 
Normal,  salts,  231 

solutions,  99 

OBSERVATION,  methods  of,  24 
Octahedron,  94 
Octane,  328 
Octaves,  law  of,  272 
Oil,  illuminating,  328 
fusel,  319 

and  salts  under  the  positive  radical. 


524 


INDEX 


Oil,  -gas,  340 

of  vitriol,  259 
Oleum,  263 
Olivine,  428 
Onyx,  346 
Opal,  346 

Open  hearth  process,  497 
Ores,  354 

Organic  chemistry,  316 
Orpiment,  462 
Orthoclase,  348 
Osmium,  508 
Osmotic  pressure,  101 

abnormal,  206 

Oxidation,  51,  52,  75,  110,  302 
Oxides,  nomenclature  of,  51 
Oxidizing  agents,  111,  192,  194,  196, 

198,  210,  212,  492 
Oxygen,  45 

chemical  properties,  48 

physical  properties,  48 

preparation,  46 
Oyster  shells,  391 
Ozone,  138,  209 

PAINT,  luminous,  401 

white,  405,  432,  459 
Palladium,  508,  509 
Paper,  manufacture  of,  332,  446 
Paraffin,  328 

series,  327 

Paris  green,  415,  465 
Paris,  plaster  of,  400 
Parke's  process,  417 
Partition,  law  of,  103 
Passive  metals,  475 
Pearl  ash,  391 
Pearls,  367 
Pentane,  327,  328 
Pentatriacontane,  327 
Pentlandite,  505 
Perchlorates,  196 
Perchloric  anhydride,  197 
Periodic  system,  271,  276 

see  Table  opposite  inside  of  rear 

cover 

Permanent  white,  405 
Permanganates,  487,  491 
Permanganic  anhydride,  491 
Peroxides,  in  restricted  sense,  212,  488 
Petroleum,  328 
Pewter,  452 
Phenol,  294 
Phenolphthalein,  242 
Philosophers'  stone,  303 
Phosgene,  326 


Phosphine,  306,  313 
Phosphonium  compounds,  307 
Phosphoric  anhydride,  308 
Phosphorus,  303 

acids  of,  309 

chemical  relations  of,  303,  314 

compared  with  nitrogen  and  sul- 
phur, 314 

halides,  307 

hydrides  of,  306 

oxides,  308 

pentachloride,  114,  178,  180,  308 
dissociation  of,  146 

pentoxide,  308 

sulphides,  314 

tetroxide,  308 

trichloride,  114,  308 

trioxide,  308 

Photochemistry,  9,  103,  305,  320 
Photography,  267,  421 
Pink-salt,  454 
Pintsch's  oil-gas,  340 
Plaster  of  Paris,  400 
Platinum,  508,  510 

compounds  of,  510 

tetrachloride,     see    Acid,     chloro- 

platinic 

Plumbates,  458 
Plumbic,  see  Lead 
Plumbic-ion,  460 
Plumbites,  458 
Polysulphides,  255 
Porcelain,  449 
Positive,  elements,  224 

radicals,  64 
Potash,  caustic,  232 
Potassium,  361 

-aluminium  sulphate,  447 

argenticyanide,  419,  421 

arsenite,  465 

aurocyanide,  423 

bisulphate,  368 

bitartrate,  369 

bromate,  197,  365 

bromide,  362 

carbonate,  366 

chlorate,  47,  111,  195,  365 

chloride,  361 

chloroplatinate,  369,  511 

chromate,  475 

cobalticyanide,  504 

cobaltinitrate,  504 

cupric  sulphate,  416 

cuprocyanide,  415 

cyanate,  336,  367 

cyanide,  335,  367 


***  Acids  are  all  listed  under  "  acid  "  and  salts  under  the  positive  radical. 


INDEX 


525 


Potassium,  dichromate,  111,  470 
ferricyanide,  502 
ferrocyanide,  501 
fluorides,  170,  362 
hydride,  361 
-hydrogen  sulphate,  368 

sulphide,  368 

tartrate,  369 
hydroxide,  363      , 
hypobromite,  197 
hypochlorite,  189 


hypophosphite,  306 
iodate,  198, 


365 

iodide,  362 

-ion,  369 

manganate,  490 

metasilicate,  344 

nitrate,  279,  292,  365 
electrolysis  of,  216 

oxide,  364 

perchlorate,  196,  365 

permanganate,  95,  110,  491,  492 
oxidation  by,  213 

peroxide,  364 

phosphate,  118,  161 

picrate,  369 

polysulphides,  368 

pyroantimoniate,  469 

sulphate,  361,  368 

sulphide,  368 

tartrate,  369 

thiocyanate,  336 

yellow  prussiate  of,  501 

-zinc  sulphate,  433 

zincate,  432 
Powder,  smokeless,  366 
Praseodymium,  443 
"  Precipitate,  white,"  438 
Precipitation,  ionic,   formulation  of, 
238 

rule  for,  394 

theory  of,  394 

Prediction  of  new  elements,  276 
Pressure,  dissociation,  391 

measurement  of  gaseous,  58 

osmotic,  101 
abnormal,  207 

partial,  60 

solution,  102 

vapor,  78,  79 

Problems,  arithmetical,  43,  60,  143 
Product,  solubility,  386 
Propane,  327,  328 
Properties,  specific,  23 

specific  chemical,  75 
Proustite,  417,  466 

*#*  Acids  are  all  listed  under  "acid 


Prussian  blue,  502 

Puddling,  497 

Pure,  chemically,  definition  of,  23 

Putrefaction,  333 

Pyrargyrite,  417,  469 

Pyrite,  248,  256,  261,  494,  501 

Pyrites,  arsenical,  462 

copper,  409 
Pyrolusite,  487 
Pyrosulphates,  263,  265 

QUALITATIVE  analysis  (cations),  439 
Quartation,  425 
Quartz,  94,  346 
Quicklime,  322,  391 
Quicksilver,  436 

RADICALS,  64,  187,  201 

free,  208 

valence  of,  70 
Radio-activity,  484 
Radium,  484 
Reaction  formulae,  83 
Realgar,  462,  466 
Red,  fire,  404 

heat,  temperature  of,  53 

lead,  488 

Reducing  agents,  75,  213,  253 
Reduction,  53,  75,  253,  302 
Refrigeration,  281 
Relations,  chemical,  of  elements,  116, 

157 

Respiration,  289 
Reversible  actions,  46,  115,  174 

completion  of,  184 
Rhodium,  508,  509 
Rhombohedron,  9 
Rinmann's  green,  433 
Roasting,  256 
Rochelle  salt,  421 
Rock  crystal,  94,  346 
Rouge,  500 
Rubidium,  369 
Ruby  copper,  409 
Rust,  iron,  498 
Rusting,  2,  4,  52 
Ruthenium,  508 
Rutile,  460 

SALAMMONIAC,  370 

Salt,  action  of  sulphuric  acid  on,  118 

-cake,  376 

common,  374 

deposits,  108,  361,  374,  428 

Glauber's,  379 

solar,  374 
and  salts  under  the  positive  radical. 


526 


INDEX 


Salting  out,  335,  387 
Saltpeter,  279 

Bengal,  292,  365 

Chili,  292 
Salts,  acid,  171,  231 

activity  of,  229 

basic,  231 

chemical  behavior  of,  187,  201 

double,  231,  416 

Epsom,  430 

hydrolysis  of,  254,  353 

ionic  substances  from,  232 

mixed,  231 

neutral,  acid  or  alkaline  reactions 

of,  353 
alkaline  reactions  of,  254 

nomenclature  of,  186 

non-ionic  formation  of,  246 

normal  or  neutral,  231 

solubilities    of,    see    Table    inside 

front  cover 
Sand,  343 
Sandstone,  348 
Saponification,  334 
Sapphire,  446 
Saturated  solution,  definition,  107 

equilibrium,  105 

making  of,  98 

theory  of,  102 

ionic  equilibrium  in,  384,  394 
Scandium,  277,  442 
Scheelite,  483 
Schlippe's  salt,  469 
Schoenite,  368 
Scientific  method,  1 
Sea-water,  108 
Selenite,  95,  400 
Selenium,  269 
Selters  water,  322 
Separation,  25 
Series,  electromotive,  245 

in  periodic  system,  273 

paraffin,  327 
Serpentine,  348,  428 
Sesquioxides,  356 
Siderite,  494 

Siemens-Martin  process,  497 
Silica,  343 

Silicates,  classification  of,  347 
Silicon,  343 

carbide,  320 

compounds  of,  344  *-. 

Silk,  448 
Silver,  417 

bromide,  419 

chloride,  419 


Silver,  chrornate,  479 

complex  compounds  of,  419 

German,  505 

hydroxide,  420 

iodide,  419 

nitrate,  420 

orthoarsenate,  465 

orthophosphate,  313,  421 

oxide,  213 

oxides  of,  419 

-plating,  421 

salts  of,  420 
Simple  substances,  liberation  of,  67 

molecular  weights  of,  137,  146,  147 
Simultaneous  actions,  194,  196 
Slag,  355 
Smalt,  504 
Smaltite,  503 
Smithsonite,  431 
Soap,  334 

salting  out  of,  335,  387 
Soapstone,  428 
Soda,  calcined,  377 

caustic,  232,  374 

crystals,  377 

washing,  377 

-water,  323 
Sodamide,  284 
Sodium,  373 

aluminate,  67,  445 

amalgam,  283,  373 

arsenite,  465 

bicarbonate,  323,  379 

carbonate,  375 
properties,  378 

chloraurate,  424 

chloride,  374,  386 

chromate,  476 

dichromate,  477 

fluosilicate,  380 

formate,  325,  330 

hydride,  373 

-hydrogen  carbonate,  323,  379 

hydroxide,  374 

hyposulphite,  379 

"  hyposulphite,"  see  Thiosiilphatt; 

iodate,  198 

-ion,  380 

metaphosphate,  312 

metasilicate,  347,  380 

metastannate,  454 
.....nitrate,  279,  292,  375 

nitrite,  299,  375 

orthophosphate,  380 

primary,  secondary  and  tertiary, 
311 


*#*  Acids  are  all  listed  under  "acid  "  and  salts  under  the  positive  radical. 


INDEX 


r,2T 


Sodium,  orthosilicate,  346 

oxide,  375 

palmitate,  335 

pentasulphide,  379 

periodate,  198 

permanganate,  491 

peroxide,  211,  374 

phosphate,  common,  380 

polysulphides,  255 

pyroantimoniate,  380,  469 

pyrophosphate,  312 

pyrosulphate,  263 

sulphate,  379 
solubility,  106 

sulphide,  hydrolysis  of,  259 

sulphite,  256 

tetraborate,  348,  350,  380 

thiosulphate,  255,  266,  379 

tungstate,  483 

zincate,  67 
Solder,  452 
Soldering,  283,  380 
Solubilities,    104;    see    Table    inside 

front  cover 
Solubility,  curves  of,  104 

independent,  103 

influence  of  temperature  on,  105 

limits  of,  97 

measurement  of,  97 

of  gases,  97,  103 

product,  386 
Solution,  96 

colloidal,  96,  500 

heat  of,  100 

in  two  solvents,  103 

of  insoluble  substances,  395  (rule 
for),  396,  413 

pressure,  102 

saturated,  98,  102,  105,  107 

scope  of  word,  96 

solid,  97 
Solutions,  freezing-points,  204 

molar,  99 

normal,  99 

saturated,  107 

ionic  equilibrium  in,  386 

solid,  97 

standard,  241 

supersaturated,  107 
Solvay  soda  process,  378 
Solvents,  immiscible,  103 
Spar,  391 

Specific  heats  of  elements,  135 
Specific  physical  properties,  23 
Specularite,  494 
Speed  of  reaction,  75,  176 


S|  >.><>(!  of  reaction,  affected  by,  cat;ily 

sis,  54 

concentration,  178 
temperature,  53,  178 

in  homogeneous  systems,  182 
Sphalerite,  431 
Spinelle,  446 
Spirit  of  hartshorn,  281 
Spirit,  wood,  318 
Sprinklers,  470 
Stability,  meaning  of,  81 
Stalactites,  392 
Standard  solutions,  241 
Stannic  bromide,  454 

chloride,  453 

-ion,  455 

nitrate,  453 

oxide,  455 

sulphide,  455 

see  Tin 

Stannites,  453,  455 
Stannous  chloride,  453 

hydroxide,  455 

-ion,  455 

nitrate,  453 

oxide,  455 

sulphate,  453 

sulphide,  455 

see  Tin 
Starch,  332 

Stassfurt  deposits,  361,  428 
Steam,  78 
Steel,  497 

Stereotype-metal,  467 
Stibine,  467 
Stibnite,  467 
Stone,  artificial,  380 
Strontianite,  404 
Strontium,  404 

carbide,  358 

compounds,  404 

-ion,  404 

peroxide,  212 
Structure,  molecular,  156 
Strychnine,  319 
Stucco,  400 

Sublimate,  corrosive,  436 
Sublimation,  308 
Substance,  compound,  definition,  21 

definition,  23,  24 

simple,  definition,  21 
Substitution,  115 
Sugar,  cane,  332 

crystalline  form,  95 

grape,  332 

refining,  319 


***  Acids  are  all  listed  under  "acid  "  and  salts  under  the  positive  radical. 


528 


INDEX 


Sulpharsenates,  469  *  • 
Sulf  hate-ion,  264 
Sulphates,  265 

acid,  see  Hydrogen  sulphates 
Sulphides,  253 

characteristics  of,  358 
solubility  of,  in  acids,  254 
Sulphites,  266 
Sulphur,  248 

amorphous,  249 
v  chemical  relations,  255 
>>  compared    with    phosphorus    and 

nitrogen,  314 
-dioxide,  256 

•)  family,  chemical  relations  of,  270 
»  flowers  of,  248 
-j  liquid,  forms  of,  249 

monochloride,  267 
**  monoclinic,  249 
^  native,  formation  of,  252 
J  oxides  of,  256 
,  oxygen  acids  of,  258 
^recovery,  377 
rhombic,  249 
I?  roll,  248 
Vtrioxide,  257 
uses  of,  250 
vapor,  density  of,  146 
Sulphuretted  hydrogen,  see  Hydrogen 

sulphide 

Sulphuric  anhydride,  257 
Sulphurous  anhydride,  256 
Sulphuryl  chloride,  257,  268 
Superphosphate  of  lime,  401 
Supersaturated  solutions,  107 
Sylvanite,  423 
Sylvite,  108,  361 
Symbols,  39 
Synthesis,  316 

TALC,  348,  428 
Tantalum,  472 
Tartar-emetic,  468 
Tellurium,  270 

atomic  weight  of,  277 
Temperature,  its  influence,  53,  178 
Temperatures,  red  heat,  etc.,  53 
Tempering,  498 
Tension,  aqueous,  60 

dissociation,  391 
Tenacity,  order  of,  351 
Tests,  66,  114 
Thallium,  442 
Thenardite,  379 
Theory,  of  ionization,  219 


Theory,  of  precipitation,  394 
Thermochemistry,  55 
Thionyl  chloride,  267 
Thomas-Gilchrist  process,  497 
Thorite,  460 
Thorium,  460 

dioxide,  338 
Tin,  451 

action  of  nitric  acid  on,  453 

bromide,  454 

chlorides,  453 

plate,  452 

-stone,  452 

sulphides,  455 

see  Stannic  and  Stannous 
Tincal,  348 
Titanium,  460 
Titration,  242 
Topaz,  95 

Transition  point,  78 
Triolein,  334,  335 
Tripalmitin,  334 
Tristearin,  334,  335 
Tungsten,  483 
Turnbull's  blue,  502 
Turpentine,  114 
Type-metal,  457 

UNIT  weights,  chemical,  36 
Units  of  measurement,  42 
Unsaturated  compounds,  257 
Unstable,  meaning  of,  81 
Uranium,  compounds,  483 

VALENCE,  68 

definitions  of,  69,  70 

multiple,  72 

of  radicals,  70 
Vanadium,  472 
Vapor,  density,  62 

pressure,  78,  79 

tension,  91 
Varec,  164 
Vaseline,  328 

Velocity,  see  Speed  of  reaction 
Venetian  red,  500 
Verdigris,  415 
Vermilion,  438 
Vinegar,  331 
Vitriol,  blue,  415 

green,  499 

nitrous,  262 

oil  of,  259 

white,  433 
Vitriols,  433 
Volume,  molar,  129 


***  Acids  are  all  listed  under  "  acid  "  and  salts  under  the  positive  radical. 


INDEX 


529 


Volumes,  relative,  of  gases,  145 
Volumetric  analysis,  242 

WASHING  soda,  377 
Water,  77 

as  catalytic  agent,  114 

chemical  properties,  80 

composition,  84 

electrolysis, 

forms  hydrates,  82 

fraction  ionized,  230 

gas,  325,  340 

hard,  324,  335,  391 

ionization  of,  230 

of  crystallization,  see  Hydrates 

physical  properties,  78 

purification,  52,  77,  210 

soda,  322 

stability  of,  81 

surface  moisture,  80 

union  with  oxides,  51,  81 

vapor  pressure  of,  78,  79 
Waters,  natural,  77 
Weights,  atomic,  36,  130 

chemical  unit,  36 

combining,  33 

gram-molecular,  128 

molecular,  127,  205 

relative,  of  gases,  61,  145 
Weldon  process,  490 
Welsbach  mantles,  338 
Whiskey,  319 


White  lead,  459 
Wither! te,  405»« 4, 
Wolfram,  483 
Wollastonite,  348,  402 
Wood,  318,  330 
Wood's  metal,  470 
Wool,  448 

Work,  methods  of,  24 
Wulfenite,  482 

X-RAY  screens,  511 
Xenon,  291 

YEAST,  333 

ZINC,  430 

arsenide,  463 

-blende,  248,  261,  431 

carbonate,  433 

chloride,  432 

dust,  431 

hydroxide,  432 

-ion,  434 

oxide,  432 

sulphate,  433 

sulphide,  433 

-white,  432 
Zincates,  432,  433 
Zincite,  431 
Zircon,  94,  348,  460 
Zirconium,  460 


***  Acids  are  all  listed  under  "  acid  "  and  salts  under  the  positive  radical. 


IT 


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